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Chm.1.1 Analyze the structure of atoms and ions. Chm.1.1.1 Analyze the structure of atoms, isotopes, and ions. Chm.1.1.2 Analyze an atom in terms of the location of electrons. Chm.1.1.3 Explain the emission of electromagnetic radiation in spectral form in terms of the Bohr model. Chm.1.1.4 Explain the process of radioactive decay using nuclear equations and half-life.

Chemistry 1 Final Exam Review

name symbol charge relative mass

location

proton p+ +1 1 in nucleus

neutron no 0 1 in nucleuselectron e- -1 ~1/2000 out of

nucleus

The atomic number (Z) of an element is the number of protons in the nucleus of an atom of that element. The number of protons determines identity of an element, as well as many of its chemical and physical properties. The sum of the protons and neutrons in the nucleus is the mass number (A) of that particular atom. XA

Z

DALTON’S ATOMIC THEORY

1) All matter is made of tiny indivisible particles called atoms.

2) Atoms of the same element are identical; those of different atoms are different.

3) Atoms of different elements combine in whole number ratios to form compounds

4) Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.

Atoms of an element that are chemically alike but differ in mass are called isotopes of the element. Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number.

Balanced nuclear equations require that both the atomic number and the mass number must be balanced. The atomic number (the number on the bottom) determines the identity of the element.

231235U He + Th9092

42

231235U He + Th9092

42

A half-life is the time required for one-half of a radioisotope’s nuclei to decay into its products. The half-life of any particular radioisotope is constant and therefore cannot be sped up.

time amount

0 Initial amount

Half-life Initial amount ÷ 2Half-life x 2 Initial amount ÷ 4

Half-life x 3 Initial amount ÷ 8Bohr’s model of the atom helped to explain spectral lines.

BOHR MODEL OF THE ATOM

1. An electron circles the nucleus only in fixed energy ranges called orbits.

2. An electron can neither gain nor lose energy inside this orbit, but could move up or down to another

There are 2 electrons in the first energy level (orbit). There are a maximum of 8 e- in the 2nd orbit,

Electron configurations represent the way electrons are arranged in atoms. Electrons enter the lowest energy first. At most there can be only 2 electrons per orbital, and they must have opposite “spins.”

Electron configuration for iron (Fe): 1s2 2s2 2p6 3s2 3p6 4s2 3d6

Noble gas configuration for iron (Fe): [Ar] 4s2 3d6

CURRENT MODEL OF THE ATOM

Atoms are composed of electrons in a cloud around a positive nucleus.

Frequency and wavelength are inversely related. When one is high, the other is low. Energy and frequency are directly related.

FUSION – 2 or more small nuclei combine to one larger nucleusFISSION – 1 large nucleus splits to form two or more smaller nuclei

Chm.1.1 Analyze the structure of atoms and ions.

1. Which idea of John Dalton is no longer considered part of the modern view of atoms? A. Atoms are extremely small.B. Atoms of the same element have identical masses. C. Atoms combine in simple whole number ratios to form compounds. D. Atoms of different elements can combine in different ratios to form different compounds.

2. Which best describes the current atomic theory? A. Atoms consist of electrons circling in definite orbits around a positive nucleus. B. Atoms are composed of electrons in a cloud around a positive nucleus. C. Atoms can easily be split, at which time they become radioactive. D. An atom’s mass is determined by the mass of its neutrons.

3. What is the nuclear composition of uranium-235? A. 92 electrons + 143 protons B. 92 protons + 143 electronsC. 143 protons + 92 neutrons D. 92 protons + 143 neutrons

4. Which best describes the relationship between subatomic particles in any neutral atom? A. The number of protons equals the number of electrons. B. The number of protons equals the number of neutrons. C. The number of neutrons equals the number of electrons. D. The number of neutrons is greater than the number of protons.

5. Consider the spectrum for the hydrogen atom. In which situation will light be produced? A. Electrons absorb energy as they move to an excited state. B. Electrons release energy as they move to an excited state. C. Electrons absorb energy as they return to the ground state. D. Electrons release energy as they return to the ground state.

6. Which statement regarding red and green visible light is correct? A. The speed of green light is greater than that of red light. B. The wavelength of green light is longer than that of red light. C. The energy of green light is lower than that of red light. D. The frequency of green light is higher than that of red light.

7. Which color of light would a hydrogen atom emit when an electron changes from the n = 5 level to the n = 2 level?

A. red B. yellow C. green D. blue

8. What energy level transition is indicated when the light emitted by a hydrogen atom has a wavelength of 103 nm? A. n = 2 to n = 1 B. n = 3 to n = 1 C. n = 4 to n = 2 D. n = 5 to n = 2

9. Consider this diagram:

Which of the three types of radiation will penetrate the paper and wood? A. alpha, beta, gamma B. alpha and beta only C. gamma only D. beta only

10. Given the electronic configuration of 1s2 2s2 2p4 , how many electrons does this element have in its outer level?A. 2 B. 4 C. 6 D. 8

11. The half-life of phosphorus-32 is 14.30 days. How many milligrams of a 20.00 mg sample of phosphorus-32 will remain after 85.80 days?

A. 3.333 mg B. 0.6250 mg C. 0.3125 mg D. 0.1563 mg


12. Identify the element that has the abbreviated electron configuration given by [Ne] 3s2.a) O b) Mn c) Ca d) Mg

13. In the figure below, what type of nuclear activity is represented?

A. fission B. fusion C. alpha emission D. beta emission

14. Which particle will complete this reaction?

A. electron B. neutron C. nucleus D. proton

15. Which will complete this equation?

A. B. C. D.

16. Nuclear decay is a random event and A. is independent of other energy influences B. can be sped up upon heatingC. can be sped up upon cooling D. can be sped up at room temperature

Chm.1.2 Understand the bonding that occurs in simple compounds in terms of bond type, strength, and properties. Chm.1.2.1 Compare (qualitatively) the relative strengths of ionic, covalent, and metallic bonds. Chm.1.2.2 Infer the type of bond and chemical formula formed between atoms. Chm.1.2.3 Compare inter- and intra- particle forces. Chm.1.2.4 Interpret the name and formula of compounds using IUPAC convention. Chm.1.2.5 Compare the properties of ionic, covalent, metallic, and network compounds.


Metals lose electrons to attain noble gas configurations. They make positive ions, cations. The electron that is removed comes from the highest energy level. Nonmetals gain electrons to attain noble gas configurations. This means they want an octet of electrons, 8 electrons. They make negative ions, anions.

As atoms bond with each other, they decrease their potential energy, thus creating more stable arrangements

Macromolecules have large numbers of atoms linked by covalent bonds in chains or sheets (such as graphite), or in 3-dimensional structures (such as diamond and quartz). Macromolecules are in your hair and fingernails. Man-made macromolecules include polymers like PVC and nylon. A network solid is a macromolecule in which the atoms are bonded covalently in a continuous network. Examples include diamond, graphite, and quartz.

A single bond is formed from the sharing of two valence electrons. Cl2 has a single bond. A double bond occurs when atoms share two pairs of electrons, 4 electrons. O2 has a double bond. A triple bond forms when atoms share three pairs of electrons, 6 electrons. N2 has a triple bond. Triple bonds are the strongest and shortest and have the highest bond energy. Bond energy

COVALENT BONDING

Covalent compounds occur between two nonmetals or a nonmetal and hydrogen. In a covalent bond, atoms share electrons and neither atom has an ionic charge. The electronegativity difference of the bonding elements is less than 1.7. When hydrogen bonds with a nonmetal the bond is covalent. Characteristics: low MP, low BP, poor electrical conductivity, and polar nature.


Metallic BondsThe bonding in metals is explained by the electron sea model, which proposes that the atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic cations. The bond that results from this shared pool of valence electrons is called a metallic bond.

Metals conduct electricity. Metals generally have extremely high melting and

boiling points. Metals are malleable (able to be hammered into

sheets). Metals are also ductile (able to be drawn into

wire). Metals have luster (are shiny).

A mixture of elements that has metallic properties is

IONIC BONDING

Anions and cations are involved in ionic bonding and are held together by opposite charges, electrostatic attraction. The bond is formed through the transfer of electrons. Ionic bonds occur between metals and nonmetals. The electronegativity difference of the bonding elements is greater than 1.7. Characteristics: high MP, high BP, brittle, and high electrical conductivity

Naming Binary Ionic Compounds

Binary ionic compounds are composed of a metal bonded with a nonmetal.

Name the metal ion using a Roman numeral in parenthesis if necessary.

Follow this name with the name of the nonmetal ion.

Examples: Ca3P2 = calcium phosphidePbCl4 = lead (IV) chloride

ACIDS

Acids will always contain one or more hydrogen ions next to an anion. The anion determines the name of the acid.

Binary acids contain hydrogen and an anion whose name ends in –ide. When naming the acid, use the prefix hydro- and change -ide to -ic acid.

Example: HCl = hydrochloric acid

The acid is a ternary acid if the anion has oxygen in it. The anion ends in -ate or -ite. Change the suffix -ate to –ic acid. Change the suffix -ite to -ous acid. The hydro- prefix is NOT used!

Example: HNO3 = nitric acid HNO2 = nitrous acid Naming Ternary Ionic Compounds

Ternary ionic compounds are composed of at least 3 elements. Name the metal ion, using a Roman numeral in parenthesis if necessary.Follow this name with the name of the polyatomic ion.

Examples: Na3PO4 = sodium phosphateCu2SO4 = copper (I) sulfate

MOLECULAR COMPOUNDS

Molecular compounds are made of 2 nonmetals. A molecular compound’s name tells you the number of atoms through the use of prefixes.

1 mono- 2 di- 3 tri- 4 tetra- 5 penta-6 hexa- 7 hepta- 8 octa- 9 nona- 10 deca-

The name will consist of two words. One exception is we don’t write mono- if there is only one of the first element. You will not need to criss-cross oxidation numbers.

Examples: SO2 = sulfur dioxide; N2O5 = dinitrogen pentoxide

Intermolecular forces are forces of attraction. They are what make solid and liquid molecular compounds possible. Solids have higher intermolecular forces than liquids, and liquids have higher intermolecular forces than gases. The three intermolecular forces are hydrogen bonds, dipole–dipole forces and London dispersion forces.

H-bond is an attraction between molecules when H is bonded to N, O, or F.

Dipole-dipole attractions occur between polar molecules.

London dispersion forces (also called van der Waals) occurs when electrons of one molecule attracted

A polar molecule has uneven charge distribution; it has a region of partial positive charge and a region of partial negative charge.

Zn2+

Ag1+

17. What is the name of the compound PbO2?A. lead oxide B. lead (II) oxide C. lead oxide (II) D. lead (IV) oxide

18. What is the name of HCl (aq)?A. chloric acid B. hydrochloric acid C. hydrogen chloride D. perchloric acid

19. What is the chemical formula for calcium nitrate?A. CaNO3 B. Ca(NO2)2 C. Ca(NO3)2 D. Ca3N2

20. Which is the correct formula for dinitrogen pentoxide?A. N4O B. NO2 C. N2O5 D. NO4

21. What type of bonding is associated with compounds that have the following characteristics?• high melting points• conduct electricity in the molten state• solutions conduct electricity• normally crystalline solids at room temperature.

A. covalent B. ionic C. hydrogen D. metallic

22. Which pair of elements would most likely bond to form a covalently bonded compound?A. sodium and fluorine B. barium and chlorineC. phosphorus and oxygen D. magnesium and sulfur

23. Based on the VSEPR theory, what is the molecular geometry of a molecule of PI3?A. linear B. tetrahedral C. trigonal planar D. trigonal pyramidal

24. Which is a unique characteristic of the bonding between metal atoms?A. Atoms require additional electrons to reach a stable octet.B. Atoms must give away electrons to reach a stable octet.C. Atoms share valence electrons only with neighboring atoms to reach a stable octet.D. Delocalized electrons move among many atoms creating a sea of electrons.


VSEPR SHAPES AND ANGLES

LINEAR 180° EITHER only 2 elements OR central element surrounded by 2 elements and NO dots

TRIGONAL PLANAR 120° central element surrounded by 3 elements and NO dots

BENT <120° central element surrounded by 2 elements and ONE pair of dots

TETRAHEDRAL 109.5° central element surrounded by 4 elements and NO dots

TRIGONAL PYRAMIDAL 107° central element surrounded by 3 elements and ONE pair of dots

BENT 104.5° central element surrounded by 2 elements and TWO PAIRS of dots

LINEAR 180°VSEPR

VSEPR stands for Valence Shell Electron Pair Repulsion. It predicts the three dimensional geometry of molecules. The electron pairs try to get as far away as possible to minimize repulsion. VSEPR is based on the number of pairs of valence electrons, both bonded and unbonded. An unbonded pair of electrons is referred to as a lone pair.

Polar and Nonpolar Covalent Bonds and Molecules

When the atoms in a bond are the same, the electrons are shared equally. This results in a nonpolar covalent bond. Diatomic elements (Br2, I2 N2, Cl2, H2, O2 and F2) have pure nonpolar covalent bonds. Nonpolar molecules are symmetrical because there are no unshared electrons around the central atom.

The bond that forms when electrons are shared unequally is called a polar covalent bond. There are either unshared electrons around the central atom or there are different elements bonded with the

25. What type of intermolecular force is found in carbon dioxide, CO2?A. hydrogen B. dipole C. London dispersion D. Metallic

26. A triple bond has ________ bond energy and is _______ in length than a double bond.A. greater, shorter B. greater, longer C. smaller, shorter D. smaller, longer

27. How many lone pair electrons should be added to one nitrogen atom in the following dash diagram for hydrazine?

A. 0 B. 1 C. 2 D. 3

Chm.1.3 Understand the physical and chemical properties of atoms based on their position on the Periodic Table. Chm.1.3.1 Classify the components of a periodic table (period, group, metal, metalloid, nonmetal, transition). Chm.1.3.2 Infer the physical properties (atomic radius, metallic and nonmetallic characteristics) of an element based

on its position on the Periodic Table. Chm.1.3.3 Infer the atomic size, reactivity, electronegativity, and ionization energy of an element from its position on

the Periodic Table.


PERIODIC TREND

Top to Bottom

Left to Right

Atomic Radius (Size)

increasedecreas

e

Ionization Energydecreas

eincrease

Electronegativitydecreas

eincrease

Metals, Nonmetals and Metalloids

Metals are elements that have luster, conduct heat and electricity, usually bend without breaking (malleable) and are ductile. Most have extremely high melting points. Reactivity increases as you go down within a group for metals. With metals the greater the tendency to lose electrons, the more reactive the metal is. Reactive metals have low ionization energies and low electronegativities.

Most nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are brittle when solid. Many are gases at room temperature; those that are solids lack the luster of metals. Their melting points tend to be lower than those of metals. Reactivity decreases as you go down within a group for nonmetals. Smaller nonmetals have greater nuclear charge because the outer electrons are closer to the nucleus. Thus the tendency to attract electrons will increase.

Metalloids have some chemical and physical properties of metals and other properties of nonmetals. In the periodic table, the metalloids lie along the border between metals and nonmetals. Metalloids include B, Si, Ge, As, Sb, Te, and At.

PERIODICITY

Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element. Nonmetals attract electrons to achieve an octet (8 valence electrons), so fluorine, the most active nonmetal, has the highest electronegativity.

Ionization energy (IE) is the amount of energy required to completely remove an electron from a gaseous atom. Nonmetals do not want their electrons removed because they gain to achieve an octet, so nonmetals have high ionization energies.

Atomic size is influenced by two factors. Energy Level – A higher energy level is further away. Charge on nucleus - More charge (protons) pulls

Ionic Size

As you might expect, the loss of electrons produces a positive ion (a cation) with a radius that is smaller than that of the parent atom. Metals form cations because they tend to lose electrons in order to achieve the stability of a filled or even half-filled octet. Conversely, when an atom gains electrons, the resulting negative ion (an anion) is larger than the parent atom. Nonmetals form anions in order to achieve the

Groups are vertical columns on the periodic table, while periods are horizontal rows. Main group elements (Group A elements) in the same column have similar properties, the same number of valence electrons, and the same oxidation number.

28. Which electron configuration represents a transition element?A. 1s22s22p3 B. 1s22s22p63s2 C. 1s22s22p63s23p64s23d7 D. 1s22s22p63s23p64s23d104p4

29. Which correctly lists four atoms from smallest to largest radii?A. I, Br, Cl, F B. F, I, Br, Cl C. Si, P, S, Cl D. Cl, S, P, Si

30. Which best explains why cations are smaller than the atoms from which they are formed?A. The metallic atom gains electrons, causing a larger effective nuclear pull.B. The metallic atom loses electrons, resulting in loss of an entire energy level.C. The nonmetallic atom gains electrons, causing a larger effective nuclear pull.D. The nonmetallic atom loses electrons, resulting in loss of an entire energy level.

31. Which have the lowest electronegativities?A. alkali metals B. halogens C. rare earth elements D. transition metals

32. The compound formed between element X and oxygen has the chemical formula X2O. Which element would X most likely represent?

A. Fe B. Zn C. Ag D. Sn

33. An element X has the outer ground state electronic configuration ns2np5. Which statement about X is incorrect? A. X forms diatomic molecules. B. X forms an X– ion.C. Compared to other elements in its period, X has a relatively low first ionization energy. D. X is in group 17.

34. Which of the following groups of elements contains metals and non-metals?


Representative Elements/ Main Group/ Group A Elements are shaded.

Transition Elements are white.

A. Group 3. B. Group 2. C. Group 14. D. Group 18.35. Which element is the odd one out because it does not have a high melting point?

A. Carbon. B. Molybdenum. C. Cesium. D. Rhenium.

36. Which list below correctly gives the relative values of ionization energy (IE1) for the elements stated?A. Ne > F > O < N > C > B B. Ne < F < O < N < C < B C. Ne > F > O > N > C > B D. Ne < F < O > N < C < B

37. Which statement about metallic radii (rmetal) is incorrect? A. Values of rmetal increase across the first row of the f-block. B. Values of rmetal increase down group 2.C. Values of rmetal increase down group 13 from Al onwards. D. Values of rmetal increase down group 1.

38. Which of the following statements about ionic radii is incorrect? A. The ionic radius of Fe3+ is less than that of Fe2+. B. The ionic radius of F– is less than that of Cl–.C. Values of the ionic radii of M2+ increase down group 2. D. The ionic radius of Br– is less than the covalent radius of Br.

39. Match the highest oxidation state available to one or more elements in a given group to the group number. Which pairing is incorrect?

A. Oxidation state +4; group 14. B. Oxidation state +1; group 1.C. Oxidation state +9; group 9. D. Oxidation state +3; group 15.

Chm.2.1 Understand the relationship among pressure, temperature, volume, and phase. Chm.2.1.1 Explain the energetic nature of phase changes. Chm.2.1.2 Explain heating and cooling curves (heat of fusion, heat of vaporization, specific heat, melting point, and

boiling point). Chm.2.1.3 Interpret the data presented in phase diagrams. Chm.2.1.4 Infer simple calorimetric calculations based on the concepts of heat lost equals heat gained and specific

heat. Chm.2.1.5 Explain the relationships among pressure, temperature, volume, and quantity of gas, both qualitative and

quantitative.


A phase diagram typically has three regions, each representing a different phase and three curves that separate each phase. The critical point indicates the critical pressure and the critical temperature above which a substance cannot exist as a liquid. The triple point is the point on a phase diagram that represents the temperature and pressure at which three phases of a substance can coexist. “Normal” means 1 atm.

Phase Change

States of matter Endo or exothermic

Kinetic energy

Potential Energy

sublimation solid → gas endo constant increasedeposition gas → solid exo constant decreasemelting solid → liquid endo constant increasefreezing liquid → solid exo constant decreasevaporization liquid→ gas endo constant increasecondensation gas → liquid exo constant decrease

Kinetic energy is energy of motion. Kinetic energy increases as the temperature of molecules goes up. Chemical systems contain both kinetic energy and potential energy.

Gases are more soluble in cooler liquids and under greater pressure.

For gases, pressure and volume are inversely proportional. As one goes up, the other will go down. Temperature is directly proportional to BOTH volume and pressure.

STP = standard temperature (273 K) and standard pressure (1 atm = 101.3 kPa = 760 mm Hg)

40. Which substance listed in the table is a liquid at 27°C?Melting Point (°C) Boiling Point (°C)

I 28 140II -10 25III 20 140IV -90 14

A. I B. II C. III D. IV

41. What happens to the pressure of a constant mass of gas at constant temperature when the volume is doubled?A. The pressure is doubled. B. The pressure remains the same.C. The pressure is reduced by ½. D. The pressure is reduced by ¼.

42. The total pressure in a closed vessel containing N2, O2 and CO2 is 30 atm. If the partial pressure of N2 is 4 atm, and the partial pressure of O2 is 6 atm, what is the partial pressure of CO2?

A. 20 atm B. 30 atm C. 40 atm D. 50 atm

43. What is the pressure, in atmospheres, exerted by a 0.100-mol sample of oxygen in a 2.00-L container at 273°C?A. 4.48 × 10−1 atm B. 2.24 × 100 atm C. 1.12 × 103 atm D. 2.24 × 103 atm

44. Consider this phase diagram:

45. At what temperature does the normal boiling point occur?A. 45°C B. 60°C C. 100°C D. 110°C

46. What happens when energy is removed from liquid water?A. Molecules slow down, and more hydrogen bonds are formed.B. Molecules slow down, and more hydrogen bonds are broken.C. Molecules move faster, and more hydrogen bonds are formed.D. Molecules move faster, and more hydrogen bonds are broken.

47. The gases helium, neon, and argon are in separate containers at 55°C. Which is true about the kinetic energy of the gases?A. Helium has the lowest mass and therefore greatest kinetic energy.B. They each have a different kinetic energy.C. Argon has the greatest mass and therefore the greatest kinetic energy.D. They all have the same average kinetic energy.

48. A piece of metal is heated in a Bunsen burner flame and then immersed in a beaker of cool water.

Which statement best describes the effect of the temperature changes on the kinetic energy of the particles?A. Kinetic energy of metal atoms decreases in the flame.B. Kinetic energy of water molecules increases when the heated metal is immersed.C. Kinetic energy of water molecules decreases when the heated metal is immersed.D. Kinetic energy of metal atoms increases when immersed in the cooler water.


49. This is a heating curve for a substance.

Between points X and Y, which would be observed?A. Solid and liquid will be present. B. Only vapor will be present.C. Liquid and vapor will be present. D. Only liquid will be present.

50. An open container of water is brought to a boil and heated until all of the water is converted to water vapor.Which describes the changes in the water molecules?

A. The molecules speed up and move farther apart. B. The molecules speed up and move closer together.C. The molecules slow down and move farther apart. D. The molecules slow down and move closer together.

51) 6.00 g of gold was heated from 20.0°C to 22.0°C. How much heat was applied to the gold?A. 1.55 J B. 15.5 J C. 17.0 J D. 32.5 J

52. A student has a beaker containing 55 g of water at 100°C. How much heat is needed to convert the water to steam?A. 120,000 J B. 18,000 J C. 2,200 J D. 330 J

53. How many grams of ice will melt at 0°C if the ice absorbs 420. J of energy?A. 0.186 g B. 0.795 g C. 1.26 g D. 5.38 × 104 g

54. An 18.0-g piece of an unidentified metal was heated from 21.5°C to 89.0°C. If 292 J of heat energy was absorbed by the metal in the heating process, what was the identity of the metal? Use the specific heat table below.

A. calcium B. copper C. iron D. silver

55. Consider this phrase diagram above. What process is occurring when a substance changes from point X (−130°C and 50 kPa) to point Y (30°C and 100 kPa)?

A. boiling B. freezing C. melting D. sublimation

56. When a chemical cold pack is activated, it becomes cool to the touch. What is happening in terms of energy?A. An exothermic reaction is occurring, absorbing cold from its surroundings.B. An exothermic reaction is occurring, releasing heat to its surroundings.C. An endothermic reaction is occurring, releasing cold to its surroundings.D. An endothermic reaction is occurring, absorbing heat from its surroundings.

57. What is the volume of two moles of hydrogen gas at STP?A. 44.8 L B. 22.4 L C. 11.2 L D. 2.00 L


Chm.2.2 Analyze chemical reactions in terms of quantities, product formation, and energy. Chm.2.2.1 Explain the energy content of a chemical reaction. Chm.2.2.2 Analyze the evidence of chemical change. Chm.2.2.3 Analyze the law of conservation of matter and how it applies to various types of chemical equations

(synthesis, decomposition, single replacement, double replacement, and combustion). Chm.2.2.4 Analyze the stoichiometric relationships inherent in a chemical reaction. Chm.2.2.5 Analyze quantitatively the composition of a substance (empirical formula, molecular formula, percent

composition, and hydrates).

TYPES OF REACTIOReactions fall into 5 categories:

1. Synthesis – only 1 product

2. Decomposition – only 1 reactant

3. Single replacement – one element switches Places with another element in a compound

4. Double replacement – elements in 2 compoundsswitch places

5. Combustion – products are carbon dioxide (CO2) and water

58. How many molecules are contained in 55.0 g of H2SO4?A. 0.561 molecule B. 3.93 molecules C. 3.38 x 1023 molecules D. 2.37 x 1024

molecules

59. How many moles are in 59.6 grams of BaSO4?A. 0.256 mole B. 3.91 moles C. 13.9 moles D. 59.6 moles


Being lower in potential energy makes reactants or products more stable. Activated complex = top of

peak Activation energy = top of

peak minus reactants ΔH = products minus

reactants + ΔH is endothermic (heat

absorbed) and −ΔH is exothermic (heat released).

The Collision Theory

Molecules must collide in order to react, and they must collide in the correct or appropriate orientation and with sufficient energy to equal or exceed the

A catalyst lowers the activation energy.

Signs a chemical reaction has occurred are based on the following criteria: Precipitate formation Gas formation Color Change (not just a change in color

intensity as a result of dilution) Temperature change

In molecular equations, the formulas of the compounds are written as though all species existed as molecules or whole units.

3NaOH + FeCl3 → Fe(OH)3 + 3NaCl

An ionic equation shows dissolved ionic compounds in terms of their free ions.

3Na+ + 3OH− + Fe3+ + 3Cl− ® Fe(OH)3 + 3Na+ + 3Cl−

The net ionic equation indicates only the species that actually take part in the reaction.

3OH− (aq) + Fe3+ (aq) ® Fe(OH)3 (s)

Product testing - A test for oxygen gas: Take a glowing splint and place it in a sample of gas, if it re-ignites the gas is oxygen. A test for hydrogen gas: Use a flaming wooden splint. H2 gas will make a squeaky pop when lit in air. Care needs to be taken with this test as large amounts of hydrogen are very explosive in air. A test for carbon dioxide gas: CO2 will extinguish a flame. A second test for carbon dioxide gas: CO2 is bubbled through limewater, and calcium carbonate solid is formed causing the limewater to become cloudy.

60. Analysis shows a compound to be, by mass, 43.8% N, 6.2% H, and 50.0% O. Which is a possible molecular formula for the substance?

A. NH4NO2 B. NH4NO3 C. NH3OH D. N2OH61. If a sample of magnesium has a mass of 60. g, how many moles of magnesium does the sample contain?

A. 1.1 moles B. 1.2 moles C. 2.0 moles D. 2.5 moles

62. What is the percent by mass of iron in the compound Fe2O3?A. 70% B. 56% C. 48% D. 30%

63. Metallic sodium reacts violently with water to form hydrogen and sodium hydroxide according to the balanced equation:

2Na + 2H2O→2NaOH + H2How many moles of hydrogen gas are generated when 4.0 moles of sodium react with excess water?

A. 1.0 mole B. 2.0 moles C. 3.0 moles D. 4.0 moles

64. A compound has an empirical formula of CH2O and a molecular mass of 180 g. What is the compound’s molecular formula?

A. C3H6O3 B. C6H12O6 C. C6H11O7 D. C12H22O11

65. Consider this reaction:3Ca (s) 2H3PO4 (aq) →Ca3(PO4)2 (s) 3H2 (g)

How many moles of calcium are required to produce 60.0 g of calcium phosphate?A. 0.145 mole B. 0.194 mole C. 0.387 mole D. 0.581 mole

66. According to the equation 2H2O (l)→2H2 (g) + O2 (g), what mass of H2O is required to yield 22.4 L of O2 at STP?

A. 12 g B. 18 g C. 24 g D. 36 g

Consider this reaction:3Mg (s) + 2H3PO4 (aq) → Mg3(PO4)2 (s) + 3H2 (g)

67. How many grams of magnesium phosphate should be produced if 5.40 grams of magnesium react with excess phosphoric acid?

A. 1.80 grams B. 19.5 grams C. 58.4 grams D. 175 grams

68. Methane (CH4) is burned in oxygen according to this balanced chemical equation:CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)

What volume of carbon dioxide is formed when 9.36 liters of methane are burned in excess oxygen at STP?

A. 9.36 L B. 15.0 L C. 18.7 L D. 22.4 L

69. Consider this reaction: NH3 (g) HCl (g) →NH4Cl (s) Which type of reaction does this equation represent?

A. combustion B. decomposition C. single replacement D. synthesis

70. Which equation represents a single replacement reaction that can occur?A. F2 2NaCl→2NaF Cl2 B. Cl2 2NaF → 2NaCl

F2C. Cu 2NaCl→ CuCl2 2Na D. Zn 2NaF→ZnF2 2Na

71. What products are formed when the metal potassium is added to water?A. K and H2O B. KOH and H2O C. K2O and H2 D. KOH and H2

72. When Na2O reacts with H2O, what is produced?A. HNaO2 B. Na H2O C. NaO H2 D. NaOH

73. Which equation is correctly balanced?A. Cu + H2SO4 → CuSO4 + H2O + SO2 B. 2Na + 2H2O →2NaOH + H2C. 2Fe + 3O2 → Fe2O3 D. 4Cu + S8 →8Cu2S


74. What coefficients are required to balance this equation?_Fe2O3 _CO →_Fe _CO2

A. 2, 6, 3, 6 B. 1, 3, 2, 3 C. 1, 1, 2, 2 D. 1, 1, 2, 1

75. Which example indicates that a chemical change has occurred?A. When two aqueous solutions are mixed, a precipitate is formed.B. As ammonium nitrate dissolves in water, it causes the temperature of the water to decrease.C. Alcohol evaporates when left in an open container.D. Water is added to blue copper (II) chloride solution. The resulting mixture is lighter blue in color.

76. An aqueous solution of silver nitrate is added to an aqueous solution of iron (II) chloride. Which is the net ionic equation for the reaction that occurs?

A. AgNO2 (aq) FeCl (aq)→AgCl (s) FeNO2 (aq)B. 2AgNO3 (aq) FeCl2 (aq)→ 2AgCl (s) Fe(NO3)2 (aq)C. 2Ag+1 (aq) NO3-1 (aq) Fe+2 (aq) Cl2 (g)→2AgCl (s)D. 2Ag+ (aq) 2Cl− (aq)→2AgCl (s)

77. This graph represents the change in energy for two laboratory trials of the same reaction.

Which factor could explain the energy difference between the trials?A. Heat was added to trial #2. B. A catalyst was added

to trial #2.C. Trial #1 was stirred. D. Trial #1 was cooled.

Chm.3.1 Understand the factors affecting rate of reaction and chemical equilibrium. Chm.3.1.1 Explain the factors that affect the rate of a reaction (temperature, concentration, particle size and presence

of a catalyst). Chm.3.1.2 Explain the conditions of a system at equilibrium. Chm.3.1.3 Infer the shift in equilibrium when a stress is applied to a chemical system (Le Chatelier’s Principle).


HOW TO INCREASE REACTION RATE

1) Use more reactive reactants2) Increase concentrations of reacting particles 3) Increase the surface area of reactants4) Increase the temperature at which a reaction

occurs5) Add a catalyst 6) Compress gas reactants

Le Châtelier’s Principle: if a system at equilibrium is disturbed, the system will move in such a way as to counteract the disturbance.• Adding a reactant or product shifts the

equilibrium away from the increase. Removing a reactant or product shifts the equilibrium towards the decrease.

• As volume is decreased, pressure increases. An increase in pressure favors the direction that has fewer moles of gas. In a reaction with the same number of product and reactant moles of gas, pressure has no effect.

• Adding heat (i.e. heating the vessel) favors away from the increase: If DH > 0 (endothermic), adding heat favors the forward reaction. If DH < 0 (exothermic), adding heat favors the reverse reaction.

• A catalyst will decrease the time taken to reach equilibrium. A catalyst does not

EQUILIBRIUM CONSTANT

Consider the Haber Process:

N2(g) + 3H2(g) 2NH3(g)We ignore the concentrations of pure liquids and pure solids in equilibrium constant expressions (K). Pure liquids and solids have a fixed concentration which will always be constant.

If K >> 1, then products dominate at equilibrium and equilibrium lies to the right. If K << 1, then reactants dominate at equilibrium and the equilibrium lies to the left.

The numerical value of K only changes with temperature.

Entropy (S) is a measure of the disorder or randomness of the particles that make up a system. A transformation (shift) will not proceed in any direction that reduces the amount of total entropy. The entropy of a solid is less than that of a liquid. The entropy of a liquid is less than that of a gas.

At equilibrium, the rate of the forward reaction is equal to the rate of the reverse reaction and the reactant and product concentrations are constant even though

Coefficients are used as exponents.

Products go on top, and reactants go on the bottom.

78. Consider this balanced chemical equation: 2H2O (aq) → 2H2O (l) + O2 (g) Which will increase the rate of the reaction?A. increasing pressure on the reaction B. decreasing concentration of the reactantsC. adding a catalyst to the reaction D. decreasing the temperature of the reaction

79. For the reaction A+ (aq) + B− (aq) → AB (s) increasing the temperature increases the rate of the reaction. Which is the best explanation for this happening?

A. The pressure increases, which in turn increases the production of products.B. The concentration of reactants increases with an increase in temperature.C. The average kinetic energy increases, so the likelihood of more effective collisions between ions increases.D. Systems are more stable at high temperatures.

80. Which statement explains why the speed of some reactions is increased when the surface area of one or all the reactants is increased?

A. increasing surface area changes the electronegativity of the reactant particlesB. increasing surface area changes the concentration of the reactant particlesC. increasing surface area changes the conductivity of reactant particlesD. increasing surface area enables more reactant particles to collide

81. Consider the equilibrium: CH3CO2H + CH3OH CH3CO2CH3 + H2O Which statement is correct?A. The equilibrium shifts to the left-hand side if CH3CO2H is added to the system.B. The equilibrium shifts to the right-hand side when extra CH3CO2CH3 is added.C. Adding additional CH3OH results in the formation of more CH3CO2CH3.D. The addition of CH3CO2H and CH3OH together has no effect on the position of the equilibrium.

82. Consider the equilibrium: Cr2O72–(aq) + 2OH–(aq) 2CrO4

2–(aq) + H2O(l) Which statement is correct?Cr2O7

2–(aq) is orange, while CrO42–(aq) is yellow.

A. Adding aqueous NaOH shifts the equilibrium to the right-hand side.B. The color of the aqueous solution is unaffected by the addition of acid.C. Adding acid shifts the equilibrium to the right-hand side.D. Adding acid has no effect on the position of the equilibrium.

83. For an irreversible chemical reaction, the rate will be affected by changes in all of these factors except:A. temperature B. presence of a catalyst C. concentration of products D. surface area of a reactant

84. When a catalyst is added to a system at equilibrium, a decrease occurs in theA. activation energy B. heat of reaction C. potential energy of the reactants D. potential energy of the products

85. Which statement describes characteristics of an endothermic reaction?A. The sign of H is positive, and the products have less potential energy than the reactants. B. The sign of H is positive, and the products have more potential energy than the reactants. C. The sign of H is negative, and the products have less potential energy than the reactants. D. The sign of H is negative, and the products have more potential energy than the reactants.

86. In which equilibrium is the forward reaction favored by decreasing the external pressure?A. 2SO2(g) + O2(g) 2SO3(g) B. 2CO(g) + O2(g) 2CO2(g)C. 2NO(g) + O2(g) 2NO2(g) D. 2NH3(g) N2(g) + 3H2(g)

87. Which conditions will increase the rate of chemical reaction?A. decreased temperature and decreased concentration of reactantsB. decreased temperature and increased concentration of reactantsC. increased temperature and decreased concentration of reactantsD. increased temperature and increased concentration of reactants

88. Given the reaction at equilibrium: N2(g) + O2(g) 2NO(g) as the concentration of N2(g) increases, the concentration of O2(g) will

A. decrease B. increase C. remains the same


89. Which is a property of a reaction that has reached equilibrium?A. The amount of products is greater than the amount of reactants. B. The amount of products is equal to the amount of reactants. C. The rate of the forward reaction is greater than the rate of the reverse reaction. D. The rate of the forward reaction is equal to the rate of the reverse reaction.

90. Given the equilibrium system at 25°C: NH4Cl(s) NH4+(aq) + Cl-(aq) (ΔH = +3.5 kcal/mol)

What change will shift the equilibrium to the right?A. decreasing the temperature to 15°C B. dissolving NaCl crystals in the equilibrium mixtureC. increasing the temperature to 35°C D. dissolving NH4NONH3 crystals in the equilibrium mixture

Chm.3.2 Understand solutions and the solution process. Chm.3.2.1 Classify substances using the hydronium and hydroxide concentrations. Chm.3.2.2 Summarize the properties of acids and bases. Chm.3.2.3 Infer the quantitative nature of a solution (molarity, dilution, and titration with a 1:1 molar ratio). Chm.3.2.4 Summarize the properties of solutions. Chm.3.2.5 Interpret solubility diagrams. Chm.3.2.6 Explain the solution process.


Gas solubility decreases with increasing temperature and decreasing pressure.

On a line means saturated (a solution which holds the maximum amount of solute per amount of the solution under the given conditions).

Below a line means unsaturated (the amount of solute dissolved is less than the maximum that could be dissolved).

Above a line means

Solutions are homogeneous mixtures in a single phase.

“Like dissolves like.” Polar solvents dissolve polar solutes.

Making SolutionsA salt dissolves faster if

it is stirred or shaken, the particles are made

smaller, the temperature is

increased.

ELECTROLYTES

An electrolye has free ions (such as Na+1 and Cl-1) in the solution and is able to conduct electricity and make a light bulb burn brightly. When sugar dissolves in water, there are no free ions to conduct electricity.

COLLIGATIVEPROPERTIES

There are 4 colligative properties. Adding a solute to a solvent does the following:

Lowers Vapor PressureElevates Boiling Point

Depresses Freezing PointIncreases Osmotic Pressure

Ultimately, colligative properties depend only on the number of dissolved

particles in solution.

An Arrhenius acid is a substance that produces hydronium ions (H3O+) when it dissolves in water. An Arrhenius base is a substance that produces hydroxide ions, OH–, when it dissolves in water. A Brønsted-Lowry acid is a proton (H+) donor. A Bronsted-Lowry base is a proton acceptor.

pH = - log [H+]

pH + pOH = 14.00

Properties of Acids

Acids taste sour. Lemon juice and vinegar, for example, are both aqueous solutions of acids. Acids conduct electricity; they are electrolytes. Acids cause litmus paper to turns red and the indicator phenolphthalein to turn clear and colorless. Acids react with active metals to form hydrogen gas. This property explains why acids corrode most metals. Acids react with

A compound is soluble due to the attraction between solvent and solute particles. Consider NaCl in water: Na+ is attracted to O2− in water, while Cl− is attracted to H+

.

91. Which will increase the solubility of most solid solutes?A. decreasing the temperature B. decreasing the amount of solvent at constant temperatureC. increasing the temperature D. increasing the amount of solute at constant temperature

92. Which 0.1 M solution will turn phenolphthalein pink?A. NaOH B. HCl C. H2O D. NaCl

93. A water sample was found to have a pH of 6 at 25°C. What is the hydroxide concentration in the water sample?A. 1 × 10−8 M B. 6 × 10−8 M C. 1 × 10−6 M D. 6 × 10−6 M

94. What is the pH of a solution of KOH with a hydroxide concentration of [OH−] = 1 × 10−4 M?A. −10 B. −4 C. 4 D. 10

95. In a titration experiment, if 30.0 mL of an HCl solution reacts with 24.6 mL of a 0.50-M NaOH solution, what is the concentration of the HCl solution?

A. 0.41 M B. 0.61 M C. 1.5 M D. 370 M

96. A student wishes to prepare approximately 100 milliliters of an aqueous solution of 6 M HCl using 12 M HCl. How many milliliters of 12 M HCl must be used?

A. 100 mL B. 200 mL C. 50 mL D. 12 mL

97. Which substance can be classified as an Arrhenius acid?A. HCl B. NaCl C. LiOH D. KOH

98. Which solution will change red litmus to blue?A. HCl B. NaCl C. CH3OH D. NaOH

99. An acidic solution could have a pH ofA. 7 B. 10 C. 3 D. 14

100. What is the pH of a 0.00001 molar HCl solution?A. 1 B. 9 C. 5 D. 4

101. As the hydrogen ion concentration of an aqueous solution increases, the hydroxide ion concentration of this solution willA. increase B. decrease C. remain the same

102. What is the total number of moles of H2SO4 needed to prepare 5.0 liters of a 2.0 M solution of H2SO4?A. 2.5 B.5.0 C. 10 D. 20


Properties of Bases

Bases taste bitter and feel slippery. Bases can be strong or weak electrolytes. They cause litmus paper to turn blue and the indicator phenolphthalein to turn pink. Bases react with acids to form water and a salt, but do not commonly react with metals.

L

mol

MDILUTIONThe number of moles of solute doesn’t change if you add more solvent.

M1 x V1 = M2 x V2

Concentration is based on the molarity of a solution. A 12 M HCl solution is said to be concentrated, while a 001 M HCl solution is said to be dilute.

The strength of a base is based on the degree of dissociation. The strength of a base does NOT depend on the molarity. A strong acid is said to dissociate completely, while a weak acid only dissociates (ionizes) slightly.

pH < 7 means acidic; H+ concentration is high

pH = 7 means neutral

pH > 7 means basic; H+ concentration is

Titration curve for the titration of a strong base with a strong acid

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