· Web viewChemistry Regents Review Concept Sheets Table of Contents Topic 1 – Math...

Libretto - Chem Review

Name____________________

Chemistry Regents Review Concept Sheets

Table of Contents

Topic 1 – Math Skills and Relationships – pg. 2

Topic 2 – Physical Behavior of Matter – pg. 3

Topic 3 – The Atom – pg. 5

Topic 4 – Nuclear Chemistry – pg. 7

Topic 5 – The Periodic Table – pg. 9

Topic 6 – Bonding – pg. 11

Topic 7 – Formulas and Equations – pg. 13

Topic 8 – The Mathematics of Formulas and Equations – pg. 15

Topic 9 – Kinetics and Equilibrium – pg. 16

Topic 10 – Properties of Solutions –pg. 18

Topic 11 – Acids, Bases, and Salts – pg. 20

Topic 12 – Oxidation-Reduction – pg. 22

Topic 13 – Organic Chemistry – pg. 24

Mr. Librettowww.elibretto.wikispaces.com

1

http://www.elibretto.wikispaces.com/


Topic #1 – Math Skills and RelationshipsMath Concepts/Skills to Master:

1. Graphing

2. Understanding Relationships (direct/indirect etc.)

3. % error = measured value−accepted valueaccepted value x 100. (See Reference Table T).

4. Significant Figures Rules:

a. Any number 1 through 9 is significantb. Any zeroes in between 1 through 9 numbers are significantc. No decimal Any zeroes at end are not significantd. Decimal Find the first number 1 through 9 all the way to the left. All

numbers from there to the right including 0 are significant. (Remember Pacific decimal point present and Atlantic decimal point absent, start counting from that side with the first nonzero number.)

Ex: 435 (3 sig figs); 4035 (4 sig figs); 403500 (4 sig figs); 4020. (4 sig figs); 4020.0 (5 sig figs); 4020.06 (6 sig figs); 0.0402 (3 sig figs); 0.4020 (4 sig figs); 0.402010 (6 sig figs).

5. K= °C + 273 (Reference Table A) Ex: 10°C = 283 K; -

273°C = 0K

6. Unit Conversions (Reference Table C):

7. Density = massvolume

8. Melting Point and Boiling Points on Reference Table S

Decimal Present

Decimal Absent

2


Topic #2 – Matter and Energy

1. Properties of Solids, liquids, gases : Solid – definite shape and volume; Liquid – definite volume, no definite shape; Gas – no definite shape or volume

2. Element – one kind of a substance – all atoms have the same atomic number (homogeneous), which cannot be decomposed. Ex: Copper (Cu)

3. Compound – two or more elements chemically combined in a definite proportion by weight (homogenous), which can be decomposed. Ex: CuCl2, H2O

4. Mixture – two or more substances physically mixed but not chemically combined (may be heterogeneous or homogenous).

Ex: NaCl(aq). Mixtures of solids are heterogeneous but solutions are homogenous.

5. Physical Change: no change in the identity of substance. Ex – melting, boiling, evaporating, freezing, condensation, sublimation. H2O(g) H2O(l)

6. Chemical Change: a substance changes into a different substance with different properties (a chemical reaction). Ex – Decomposition of water (2H2O 2H2 + O2).

7. Joule Problems: Know formula: Q=mcΔT. m= mass in grams; ΔT = change in temp.; c = specific heat

8. Joule: 4.18 J changes the temp. of 1 gram of water by 1°C. KJ = 1000 J9. Temperature: a measure of the average kinetic energy of the molecules of a

substance. Temperature scales: Kelvin (absolute) and Celcius. °K = °C + 273 but 1°K = 1°C. 10°C has higher kinetic energy than 5°C.

10. Fixed Point on Thermometer: 0°C (273 K) = melting point/freezing point of H2O and 100°C (373 K) = boiling point/condensation point of H2O. Absolute zero = 0 K (-273°C).

11. Gas Laws Problems: P1V1/T1 = P2V2/T2. a) Drop what is constant. b) Temperature must be in K!

12. Boyle’s Law: (constant Temp.) P and V vary inversely. P1V1 = P2V2 13. Charles’s Law: (constant P) V and T vary directly V1/T1 = V2/T2.

Solid Liquid Gas

3


14. Activation Energy: minimum amount of energy needed to start a reaction (all reactions need activation energy).

15. S.T.P. : standard temp. and pressure = 101.3 kPa (1 atm) and 0°C (273 K). (Table A.)

16. Density: mass/volume. 17. Phase Change Diagrams: during a phase change (ex – melting,

vaporization, freezing) the temp. stays constant. No change in K.E., only P.E.Heating Curve Cooling Curve

Endothermic: sl, lg (forward phase change)Exothermic: gl, ls (reverse phase change)

18. Kinetic Molecular Theory: Ideal Gas – no attractive forces between molecules, molecules have no volume. Real Gas – (ex. H2, O2) has attractive forces, have volume.

19. Ideal and Real gases most alike (deviate the least) when attractive forces are the weakest (need molecules apart to have the weakest forces). P.L.I.G.H.T – low pressure, ideal gas, high temperature.

20. The Smallest molecule is the most ideal – He, and other monatomic Group 18 molecules. Next most like ideal are diatomic molecules (BrINClHOF)

21. Heat of Fusion : # of joules required to melt one gram of solid at its melting point. H2O = 334 J/g (Table B). q = mHf

22. Heat of Vaporization: # of joules required to vaporize one gram of liquid at its boiling point. H2O = 2,260 J/g (Table B). q = mHv.

23. Boiling Point: the temperature at which the vapor pressure of a substance equals atmospheric pressure. (Normal B.P. = B.P corresponding to the air pressure = 1atm, 101.3 kPa).

4


24. Vapor Pressure: dependent on: a) temp. of the liquid – the higher the temp. the higher the v.p.; b) strength of the intermolecular forces of attraction – the stronger the attractive forces, the lower the v.p (Ref Table H).

25. Substances that evaporate readily have high v.p. and low b.p. Sublimation – weak attractive forces, high vapor pressure. Higher the b.p., the stronger the IMAF are.

5


Topic #3 – The Atom1. Fundamental particles of the atom:

Proton = +1 charge, mass = 1 a.m.uElectron = -1 charge, mass = ~0 (1/1836 of a.m.u) Neutron = 0 charge, mass = 1 a.m.u.

1 a.m.u. = 1 atomic mass unit = 1/12 mass if Carbon-12.

2. Nucleons : protons and neutrons (the particles in the nucleus).3. Nucleus: contains most of the mass of the atom, and has a positive charge.

Charge of the nucleus = the number of protons (called the nuclear charge). Ex: if an atom has 3 protons and 4 neutrons the nuclear charge = +3; an atom with 7 protons and 5 neutrons the nuclear charge = +7.

Charge Mass # Location

Proton +1 1 NucleusNeutro

n 0 1 NucleusElectro

n -1 0 Outside Nucleus

4. Atoms are neutral : # protons = # electrons. Ex = If an atom has 14 protons, how many electrons will it have? Answer = 14.

5. Atomic Number = # protons in the nucleus. All atoms of the same element have the same number of protons. This identifies the element. Ex: All atoms of sodium must have 11 protons.

6. Mass Number = the # of protons and the # of neutrons in the nucleus (a whole # not found on the Periodic Table).

7. Nuclear Notation: shows Mass Number = p+n; Atomic Number = p as: p+npXEx. You are given 15

8 C, how many neutrons are there? Answer = top # - bottom # = 15 – 8 = 7

8. Isotopes are atoms of the same element (same # of protons) but a different number of neutrons. Isotopes of the same element have the same number of protons and electrons. Ex: = 12

6 C = 6 protons, 6 neutrons, and 6 electrons but 146 C = 6 protons, 8 neutrons, and 6 electrons.

6


9. Average atomic mass of an element (on periodic table) is determined by taking the weighted average mass of the naturally occurring isotopes of the element. Because of this, the atomic mass listed on the periodic table is a decimal.

Average Atomic Mass = (% isotope1)(mass isotope1) + (% isotope2)(mass isotope2)…etc.

Ex: 20X = 80%, 22X = 20% therefore avg. atomic mass = (.80)(20)+(.20)(22) = 20.4 amu

10. Ernest Rutherford – gold foil experiment shows a) atoms are mostly empty space;b) atoms have a positively charged, dense center.

11. Bohr’s Model of the Atom: an electron in the ground state can absorb energy and jump to higher energy levels or an excited state; the energy is then released as wavelengths of light as the electrons fall back to the ground state. The wavelengths of light emitted are unique to each element, so an element’s identity can be found by studying the wavelength (color) of light it gives off. This is the emission spectra of light.

12. Valence Electrons: #e in the outermost principle energy level. Ex. ❑9 F has 9 total electrons, and 7 valence electrons.

13. Electron Dot Diagram: uses dots to show the number of valence electrons. Ex. Fluorine ❑9 F = 2-7 Carbon ❑6 C = 2-4

7


14. Know the following chart:

Principal Energy Levels

(P.E.L.)Number of e- in

the P.E.L.n=1 2

n = 2 8n = 3 18n= 4 32

The maximum number of electrons in each energy level = 2n2

8


Topic #4 – Nuclear Chemistry

1. Radioisotope = an isotope that is radioactive and unstable.2. Stability of nuclei – nuclei are composed of protons and

neutrons. The ratio of protons to neutrons determines the stability of the atom.

a. If the ration of protons to neutrons is not stable, the isotope is radioactive (radioisotope).

b. Any element with an atomic number great than 83 (83 or more protons) is radioactive.

3. Transmutation: different elements from reactant side to product side. Ex: 226

88 Ra 22288 Rn + 4

2 He4. Chemical Reaction: Has the same elements on both sides.

Ex: 2H2 + O2 H2O5. Artificial Transmutation: Bombard a nucleus with high-energy particles

that change it from one element to another. Transmutation reactions have two things on reactant side, and the elements change from reactants to products. Ex: 14

7 N + 42 He 17

8 O + 11H

6. Natural Transmutation: (one reactant on left side, elements change from left side to right side):

a. Alpha, Beta, positron, gamma decay (see pg. 217 of review book or Reference Table O).

b. Alpha Decay:

Atomic # decreases by 2, Mass # decreases by 4.

c. Beta Decay:23290 Th 0

−1e + 23291 Pa; Mass # remains the same; atomic number

increases by 2.d. Positron Decay:

3791K 0

+1e + 3718Ar; Mass # remains the same; atomic number decreases

by 1.9


e. Gamma Decay:33He 3

2He + 00ϒ ; no change in mass # or atomic # - high-energy

photon released.Reference Table N has decay modes.

7. Nuclear Fission: splitting of an atom to produce energy (exothermic). 23592U + 1

0n 8735Br + 146

57 La + 310n

8. Nuclear Fusion: joining light molecule together to produce energy.21H + 2

1H 42 He; difficult to initiate because positive nuclei repel each other.

9. Half-Life: the amount of time it takes for a radioactive sample to half itsradioactive mass (Reference Table N)

a. Half-life not affected by anything, including temperature or pressure.b. Reference Table T Equations – Radioactive Decay:

Fraction Remaining = 12

tT ; t = total time; T = half-life.

Number of half-life periods = tT ; t = total time; T = half-life. (Use when given two different times).

1 (if fraction) Half-Life TimeFraction

or orMass in kg Final Mass (kg)

10. Uses of radioisotopes: - used to trace the path of a chemical reaction (tracers)

a. C-14: used for dating organic substances (anything that was once living)

b. I-131: Thyroid diagnosisc. Co-60: treating cancerd. U-230: used in geological dating e. Tc-99: used to trace cancer cells

10


Topic #5 – The Periodic Table1. The arrangement of the Periodic Table is based on atomic number.2. The Periodic Law states that the chemical properties of elements are

periodic functions of their atomic numbers.3. Elements are classified as (1) metals (2) non-metals (3) metalloids. More

than 2/3 of all the elements are metals. Trends:

4. Metalloids: have properties of both metals and non-metals.5. Metals : lose electrons to form positive ions, and become smaller. Ex:

Na+ is smaller then Na0. The ion will generally have the electron structure of an inert gas. (Metals are solid at room temperature except Hg = liquid).

6. Non-metals: - gain electrons to form negative ions, and become larger. Ex: Cl- is larger than Cl0 The ion will generally have the electron structure of an inert gas. The only liquid non-metal is Bromine (Br2).

7. Group 1 metals: called alkali metals (form strongest bases); so active only exist in compounds.

8. Group 2 metals: called alkaline earth metals; less active than alkali metals.

9. Group 18 nonmetals: called Inert (noble or rare) Gases. 10. Group 16, Group 15 and Group 14: contain both nonmetals and

metalloids.11. Group 17 non metals: called halogens. This group exhibits all three

phases of matter at room temperature (F2 and Cl2 are gases, Br2 is a liquid, I2 is a solid).

12. Groups 3-12 are the transition metals. 13. Elements found in the same period have the same number of energy

levels. 14. Elements found in the same group have the same number of valence

electrons and therefore similar chemical properties. 15. The most active metals are in the lower left corner of The Periodic

Table (excluding Li) and the most active non-metals are in the upper right corner of the Periodic Table (excluding inert gases).

16. The most active elements form the most stable compounds: Ex. Group 1 Rb and Group 17, F = RbF. Very Stable.

11


17. Monatomic molecules – contain 1 atom per molecule = Group 18: He, Ne, Ar, Kr, Xe, Rn (generally nonreactive)Allotropes – different forms of the same element ex – O3, S8

18. Diatomic Molecules – 2 atoms per molecule = Br2, I2, N2, Cl2, H2, O2, F2 (BrINClHOF).

19. Transition Elements – (1) Group B elements and Group VIII can have electrons from two outermost shells involved in a reaction. (2) Form colored ions in compounds or solutions (3) have multiple oxidation states (4) Remember Cu2+ is blue, thus CuSO4 is blue in solution.

20. Xe and Kr, although expected to be inert (like the other Group O elements), can form compounds with F and O under special conditions. Ex: XeF4, XeF6 exist.

21. Van der Waals forces increase down a group (due to increasing molecular size), and thus b.p. and m.p. increases.Examples:

As Van der Waals IncreaseF.P and B.P. Increase

The smallest molecule has the lowest melting point

22. Formulas: Ex: X2O3 X is in Group 3, has a +3 charge (oxidation number)

XD X is in Group 2, has a +2 charge (oxidation number)23. States of elements: Metals are all solids at room temperature except

Hg is a liquid. Nonmetals that are gases H2, N2, O2, F2, Cl2 and Group 18. Liquid = Br2; Solids = C, P, S, I2 (nonmetals)

24. Atomic radius decreases across period due to increase in # of protons (greater nuclear charge) in nucleus (more pull with same # of energy levels).Atomic radius increases down each group due to the addition of an extra shell for each successive element (increases size substantially).

25. Carefully examine charts – Periodic Table, and Table S for atomic radius, ionization energy and electronegativity trends by element

F2

Cl2

Br2

I2

HeNeArKrXeRn

12


.

13


Topic #6 –Bo nding 1. When a bond is made energy is released (exothermic). When a bond is

broken, energy is absorbed (endothermic). Remember “Karate Chop.”2. Atoms bond together to become more stable (usually to get a stop octet

ending).3. Metals tend to lose electrons to form positively charged ions (the metal

ion is smaller than the atom), and have low electronegativity values. Ex: Na+ is smaller than Na0.

4. Non-metals tend to gain electrons to form negatively charged ions (the ion is larger than the atom), and have high electronegativity values (F is the highest) Ex: Cl- is larger than Cl0.

5. A chemical bond: results from the simultaneous attraction of electrons to two nuclei. (Usually between two non-metals.)

6. Ionic Bonds : formed between metal and non-metal atoms; created by a transfer of electrons. Ex: NaCl; MgBr2. (MINT – Metal Ionic Nonmetal Transfer of e- from metal to nonmetal).

7. Covalent Bond: formed by the sharing of two electrons between two nuclei. (Usually between two non-metals.) Ex. CO2, NH3.

8. Electronegativity: the ability of an atom to attract the electron in a bond. See Reference Table S for electronegativities. Electronegativity based on a scale of 0.0 - 4.0 By subtracting the electronegativity values of the two elements, the electronegativity difference (END) is calculated, and can be used to determine the type of bond:

Ionic Bond –electrons are transferred (metal and a nonmetal) Polar Covalent Bond – electrons are shared unequally – one element

has partial negative charge, the other a partial positive charge Non-Polar Covalent Bond – electrons are shared equally – diatomic

molecules Large END = more ionic, less covalent Small END = more covalent, less ionic

Ex: NaCl – Electronegativity difference = 3.2-0.9 = 2.3. This is IONIC MgCl2 = 3.2-1.3 = 1.9. This is IONIC. H2O = 3.4- 2.2 = 1.2 this is

Polar Covalent. Diatomic Molecules (BrINClHOF) have non-polar covalent bonding (END=O). Ex. N2 has a triple bond

9. In electron dot digrams of covalent compounds all atoms need 8 electrons around them, except for H, which has 2 electrons around it.

10. A molecule is defined as a particle, which has covalent bonds, and is the smallest unit that shows the properties of the substance. Ex: H2; H2O; CO2.

11. Polyatomic Compounds have both covalent and ionic bonds:Ex: Na2SO4 has an ionic bond between N-SO4; and a covalent bond between the S-O4.

12. K+1 and Cl-1 have the same number of electrons, and the same electron configuration.

K+1 = 18 electrons Cl- = 18 electrons

14


13.Ionization Energy: the amount of energy required to remove the most loosely bound electron. Ionization Energies are list on Reference Table S. Ex: Ionization energy of Li = 520 kJ/mol, and F=1681kj/mol; this means it takes less energy to take an electron from Li than from F.

14.Ionic Solids: Ex - NaCl, K2O, etc. a. Have high melting points, high boiling pointsb. Are hardc. Do not conduct electricity as solids, but do conduct

when dissolved in water, melted, or evaporated.15.Metallic Solids: Ex – Ag, Zn, etc.

a. Mobile electrons (“sea of electrons”)b. Conductors in solid phasec. Malleable, ductile (bendable, made into wires)d. Liquid metal at room temp = Hg

16.Molecular Solids: Ex – H2O, CH4, etc. a. Held together by weak attractive forces (Van der Waals)b. Have low melting point, low boiling pointc. Are softd. Are poor conductors

17.Network Solids: Ex – SiO2, SiC, diamond, graphitea. Held together by strong covalent bonds b. Often made of group 14 elementsc. High melting point, high boiling pointd. Are electrical insulators

18.Van der Waals Forces: exist between non-polar molecules – all diatomic, monatomic, and molecules that are not dipoles. Ex: CO2, CH4, etc.

a. Van der Waals Forces depend on the size of molecules and distance between molecules.

b. Larger molecules = stronger Van der Waals Forces.c. Closer the molecules = stronger Van der Waals Forces.d. Ex: C17H36 is a liquid at room temp, by CH4 is a gas. The larger

molecule has stronger Van der Waals Forces, which requires a higher temp. to turn it into a gas.

19.Hydrogen Bonding: between molecules, which contain Hydrogen bonded to an atom of small radius and high electronegativity (HFON). Strongest hydrogen bonds between molecules of H2O, HF, NH3. Hydrogen bonds responsible for high b.p. of H2O, HF, and NH3.

20.Polar Molecule: a dipole – polar bonds in an asymmetrical molecule shape; ends have a partial charge. Polar Molecules:

21.Non-Polar Molecule: molecule has no charged ends – molecule is symmetrical.

15


16


Topic #7 –Formulas and Equations1. Reference Table S – has the names and symbols of the elements on the

Periodic Table.2. Diatomic Molecules: (BrINClHOF) – elements that pair up when not

combined with other elements – Br2, I2, Cl2, H2, O2, F2.3. Polyatomic Ions: See Reference Table E – ions made of more than one

element and act as a unit in compounds.4. Binary Compounds – have only two elements in the formula; Polyatomic

Compounds: have more than two elements in the formula5. Formula Writing:

a. Criss-cross oxidation numbers to write subscriptsb. Any 1’s for subscript don’t get written. Ex: KCl; MgCl2c. Keep polyatomic ions in parentheses unless

there is a “1” outside it. Ex1: Mg(NO3)2 needs parentheses, telling you there are 2 Nitrogens and 6 Oxygens.Ex2: NaNO3 does not need parentheses

d. If “ate”, “ite”, or ammonium are in the name use Table E. Otherwise it will be a binary compound. Exceptions of “-ide” include peroxide, hydroxide, and cyanide. Be careful of chlorine polyatomics; there are four very similar ones.

e. Examples of compounds with polyatomic ions from Table E: Ammonium sulfite (NH4)2SO3; Magnesium Nitrate Mg(NO3)2; Aluminum hypochlorite Al(ClO)3. Examples of binary compounds: Hydrogen Peroxide H2O2; Barium Fluoride BaF.

6. Formulas and The Stock System:a. When Roman numerals are given in the name of the compound, it is the

oxidation number of the first element. b. Roman Numerals: I (+1), II (+2), III (+3), IV (+4), V (+5), VI (+6).

Ex: Copper (I) sulfide = Cu2S; Iron (III) chloride = FeCl3; Tin (IV) sulfate Sn(SO4)2; Lead (II) nitrate Pb(NO3)2

c. Naming Compounds: Make sure the first element has only one oxidation number when naming without the stock system (roman numerals)

d. Naming Binary Compounds: end in –ide. Ex: MgCl2 is Magnesium Chloride, no roman numeral since Mg can only have a +2 oxidation number.

e. Naming Polyatomic Compounds: keep the suffix of the polyatomic ion if it is second. Ex: NH4NO3 ammonium nitrate.

f. Peroxides: the subscripts are not reduced. Ex: Hydrogen Peroxide H2O2; Sodium Peroxide Na2O2

7. Naming Compounds with the Stock System:a. Solve for the oxidation number of the first element. FeI2 – oxidation

number of Fe could be +2 or +3, so the number must be written in the 17


name = Iron (II) Iodide. Ex: CuSO4 = Copper (II) sulfate; Sn3(PO4)2 = Tin (II) Phosphate; SrBr2 = Strontium Bromide (No stock necessary since Sr has only one oxidation number; no stock number ever needed for second part of compound.)

8. Types of Reactions:a. Synthesis (A + B AB): one product only.

Ex: 2H2(g) + O2(g) 2H2O (g) 4Fe(s)+ 3O2(g)2Fe2O3(s)

b. Decomposition (AB A + B): one reactant onlyEx: H2O(l) H2(g) + O2(g) or CaCO3(s)CaO(s) + CO2(g)

c. Single Replacement (A+BC AC + B): 1 element and 1 compound on both sides. Ex: Cu(s)+2AgNO3 Cu(NO3)2(aq) + 2Ag(s) or Mg(s) + 2HCl(aq) H2(g) + MgCl2 (s)

d. Double Replacement Reaction (AB + CD AD + CB): 2 compounds on both sides. Ex: AgNO3 + NaCl AgCl + NaNO3

e. Neutralization Reaction: Acid + Base Salt + Water18


Ex: H2SO4 + NaOH Na2SO4 + H2O(Write H2O as HOH) and balance:

H2SO4 + 2NaOH Na2SO4 + 2HOH9. Balancing Equations: number of atoms for each element must be the same

on both sides.a. Fe + Al2O3 Al + FeO. Balance single elements first: 9Fe +

4Al2O3 8Al + 3Fe3O4b. Keep polyatomic ions together when balancing:

Al2(SO4)3 + ZnCl2 AlCl3 + ZnSO4Balanced:

Al2(SO4)3 + 3ZnCl2 2AlCl3 + 3ZnSO4

19


Topic #8 –The Mathematics of Formulas and Equations1. Formula Mass: use the atomic mass of each element and round off to the

nearest whole number. Ex: Mg(NO3)2 = 148 g/mol

Element

# Atoms

Atomic

Mass

Formula Mass

Mg 1 x 24 = 24N 2 x 14 = 28O 6 x 16 = 96

=1482. Percent Composition = ( partwhole

) *100% Find % composition of N in Mg(NO3)2 =

28148 x 100 =19%

Find % composition of H2O in CuSO4 5H2O = 90

250x 100 = 36%. Mole Calculations (See reference Table T): number of moles =

GivenMass(g)gram−formulamass

Find the mass of 2.5 moles of Mg(NO3)2. 2.5 moles = x148g /mol . x = 370g.

3. Molecular Formula: shows the actual number of atoms in the molecule. Ex. C6H12O6.

4. Empirical Formula: show the simplest whole number ration of atoms. Ex. The empirical formula of C6H12O6 is CH2O (cannot be divided anymore and have a whole number for each).

5. Finding Molecular Formulas from Empirical Formulas: Gram-molecular weight = whole #. Then multiply subscripts of empirical

formula by that # to get molecular form. Ex. A compound has a molecular mass of 180 a.m.u. and an empirical formula CH2O. What is the molecular formula: ANS:

Formula Mass = 12 + 2 + 16 = 30. 180

30 = 6x’s

C1x6H2x6O1x6 = C6H12O66. Mole Relations in Balanced Equations

Ex: Find the number of moles of oxygen produced when 1.5 moles of KClO3 decomposes given the following reaction: 2KClO3 2KCl + 3 O2. 1.5moles2moles

= xmoles3moles

1.5KClO3 X O2 x = 2.25 moles of O2

7. Density = massvolume . Density of elements is listed on Reference Table S.

Ex: CuSO4 5H2O = 250 g/molEleme

nt#

Atoms

X Atomi

c Mass

Formula Mass

Cu 1 x 64 = 64S 1 x 32 = 32O 4 x 16 = 64

H2O 5 x 18 = 90= 250

C: 1 x 12 = 12H: 2 x 1 = 2O: 1 x 16 =

20


8. Avogadro’s Hypothesis: Equal volumes of different gases at the same temperature and pressure contain the same number of molecules.

21


Topic #9 – Kinetics and Equilibrium1. Bond-Breaking is endothermic (energy must be absorbed to break a bond).

When a bond is made, energy is released – exothermic (karate chop).2. The heat of reaction (ΔH) is the difference between the potential energy of

products and the potential energy of reactants.3.

Endothermic ExothermicEnergy is absorbed Energy is released (liberated)

Reactants have less energy than products

Products have less energy than reactants

ΔH = +kJ ΔH = -kJForm less stable products in comparison

to reactantsForm more stable products than

reactants

4. Exothermic reactions are self-sustaining, because the reaction releases enough energy to keep it going. Endothermic reactions are not self-sustaining therefore they continually need added energy to keep it going.

5. ΔH (heat of reaction) is measured in kJ (kilojoules). The ΔH values are expressed for the compounds formed, and ΔH values for common reactions are found on Table I. To find the ΔH for the reverse reaction of an equation listed, reverse the sign of the ΔH value.

6. Exothermic Reactions: A+BC+D ΔH = -kJ

ORA+BC+D+kcal

7. Endothermic Reactions: A+BC+D ΔH = +kJ

ORA+B+kJC+D

8. A –ΔH tells you the reaction is exothermic in that direction, if you switch the reaction, the opposite direction will be endothermic and have a +ΔH.Ex: For N2 +3H2 2NH3, the ΔH=-91.8kJ and is exothermic. For the reverse, 2NH3 N2 +3H2 the ΔH = +91.8kJ and is endothermic. Ex: How many kJ are required to decompose 1 mole of NH3? Ans: +91.8kJ/2 moles NH3 = 45.9kJ.

If you reverse the reaction, reverse the PE diagram

22


9. Factors Affecting Reaction Rate (speed): (Effective Molecular Collisions)a. Catalyst: speeds up a reaction by lowering the reaction energy required to

start the reaction. b. Increasing Concentration of one or more reactants increases the

reaction rate due to an increasing number of collisions.c. Increasing Temperature: increases the rate of all reactions by increasing

the number of collisions and the effectiveness of the collisions.d. For gases only: Increasing the pressure will increase the reaction rate by

increasing the number of collisions (molecules are closer together).e. Increase surface area (solids): break solids into smaller pieces to

increase the reaction rate.10. Entropy( ΔS): the amount of disorder, randomness, or lack of organization of

a system.a. Solids have the least entropy, gases the most.b. Compounds have less entropy than free elements. Ex. H2O(g) has less

entropy than H2(g) and O2(g).c. A compound dissolved in water increases in entropy. Ex. C6H12O6(s) + H2O

C6H12O6(aq).11. Exothermic Reactions (-ΔH) are favored in nature; Reactions with an increase

in Entropy (+ΔS) are favored in nature.12. Equilibrium shifts : Shifting the equilibrium point to the left makes more

reactants, shifting to the right makes more products. Think of a seesaw; the direction the reaction shifts is the side you push on to make the seesaw level again. Example:

C3H8(g)+5O2(g) 3CO2(g)+4H2O(g)+2219 kJ

a. If O2 is added, then the left side of the seesaw goes down; push on the right side to make it level = reaction shift right, which causes the CO2

and H2O to increase, and the C3H8 to decrease.

b. Pressure: an increase in pressure will cause the reaction to shift to the side with fewer gas molecules (add coefficients of gas on each side). Ex: increase in pressure will shift the reaction to the left. A decrease in pressure would do the opposite. Increase in pressure:

23


c. Temperature: An increase in temperature shift equilibrium away from the kJ, a decrease towards the kJ. Ex: The reaction in exothermic, so increasing the temperature would shift the reaction to the left.

24


Topic #10 – Properties of Solutions 1. Solution – a homogenous mixture. Ex – NaCl(aq).2. Properties of Solutions:

Are clear and do not disperse light Can have a color (transition elements) Will not settle on standing Will pass through a filter

3. Rate of Solution (How Quickly it Dissolves): Decrease the size by crushing to increase the surface area Stir Increase the temperature

4. “Likes Dissolve Likes” Polar solvents dissolve polar solutes. Ex – NaCl in water. Nonpolar solvents dissolve nonpolar solutes (oil paint in turpentine)

5. Temperature Change When dissolving, the reaction is endothermic if the water gets colder When dissolving, the reaction is exothermic if the water gets warmer.

6. Solubility – maximum quantity of a solute that can be dissolved in a certain amount of solvent or solution at a specific temperature.

7. Factors Affecting Solubility: Nature of substance (Type of substance) Temperature - solids increase with higher temp; gases

decreases with higher temp. (Ex - CO2 comes out of soda in fridge)

Pressure – only affects gases; solubility increases when pressure increases. (Ex – put cap on soda).

8. Unsaturated Solution – holds less solute than maximum. This means more can still be dissolved.

9. Saturated Solution (at equilibrium) – dissolves maximum amount; cannot dissolve any more solute without changing another factor (like temp, etc.)

10. Supersaturated Solution – temporary state that is dissolving more solute than it should; very likely will precipitate out.

11. Concentrated – holds large amount of solute relative to the amount of solvent

12. Dilute – holds small amount of solute relative to the amount of solvent.

13. Solubility Curves: See Reference Table G14. Units of Concentration: (See Ref. Table T)

a. Molarity = moles of SoluteLitersof Solvent ; b. # moles = Molarity x Volume in L c. #grams = Molarity x Volume in L x GMWEx: What is the molarity of a solution containing 82.0 g of Ca(NO3)2 in 2.0 L of solution:

25


Ans: GMW of Ca(NO3)2 = 164 g/mol 0.5 moles. M=molesLiters = 0.5moles2.0L = .250

M15. Colligative Properties: Solutes added to water will raise the boiling point

and lowers the freezing point of water. Same strength electrolytes (acid, base, or salt) will change the b.p and f.p. more than a nonelectrolyte (such as organic substances or alcohol).

a. A combination of ions and concentration determines how much the f.p. decreases and the b.p increases.

Ex: Freezing Point Highest to Lowest: 1M C6H12O6, 2M C6H12O6, 1M Mg(NO3)2, 2M Mg(NO3)2; Boiling Point Lowest to Highest: 1M C6H12O6, 2M C6H12O6, 1M Mg(NO3)2, 2M Mg(NO3)2.

Reason: Mg(NO3)2 breaks un into three ions: Mg2+ and 2 NO3-. **Count subscripts of ions and multiply by molarity. **C6H12O6 does not break into ions (covalently bonded – other examples: CH4, O2, C6H8O7, etc.)

16. Parts Per Million (ppm) = on reference Table T. ppm = gramsof solutegrams of solution x 1,000,000. (Grams of solution = grams of solute + grams of solvent).

Ex: Approximately 0.0043g of oxygen can be dissolved in 100 g of water. Express the concentration in ppm.

Ans: ppm = 0.0043g100.0043gx 1,000,000 = 43ppm

26


27


Topic #11 – Acids, Bases, Salts1. Electrolytes: A compound when melted, vaporized, or dissolved in H2O will

conduct electricity (form ions). Ex: Acids, Bases, Salts.2. Non-electrolytes: will not conduct electricity under the above conditions (does

not form ions). Ex – Organic Compounds (Except CH3COOH)3. Strong Electrolyte: strong acids, soluble bases and salts (See Table F for

solubility – nitrates vs. carbonates).4. Acids: begin with H. Ex: HCl, HC2H3O2, H2SO4, CH3COOH (acetic acid is an

exception), See Ref. Table K for a list. Form H+ (same thing as H3O+) ions.Bases: end on –OH. Ex: NaOH, NH4OH, Ca(OH)2, etc. All bases dissociate except for NH3(aq) which ionizes. Bases increase the [OH-].Salts: ionic compounds that do not begin with H or end in –OH. Ex: NaCl, MgS, MgCO3Organic: begin with C. Ex: C6H12O6 etc. Exceptions include CH3COOH (acid). Alcohols are a type of organic compounds that start with C and end in OH (like CH3OH).

5. Properties of Acids: a. Turn blue litmus red (colorless in phenolphthalein)b. Have a pH less than 7c. React with metals above H2 on Table J to form

salt and H2 gasd. React with bases to form a salt and water

(neutralization).e. Taste sourf. Conduct electricity in relation to the degree of

their ionization (more soluble = more ions = better conductor).

g. Acidic solutions contain more H+ (H3O+) ions than OH- ions.

6. Properties of Bases: a. Turn red litmus blue, pink in phenolphthaleinb. Have a pH greater than 7c. React with acids to form a salt and water (neutralization)d. Taste bitter, and feel slipperye. Conduct electricity in relation to their solubility (more soluble = more

ions = better conductor).f. Basic solution contain more OH- than H+.

7. pH scale: Compares the [H3O+] to [OH-]. H3O+ is called the hydronium ion and OH- is called the hydroxide ion.

[H3O+] x [OH-] = 10-14 (a constant.)Use this expression to find either [H If the [H3O+] increases, the [OH-] decreases, a solution becomes acidic and pH < 7. If the [H3O+] decreases, the [OH-] increases, a solution becomes basic and pH > 7. In water the [H3O+] = [OH-], and the pH = 7 which is neutral. Adding an acid decreases the pH, adding a base increases the pH.

28


8. Neutralization: a. Acid + Base salt + waterb. H+ + OH- H2O orc. H3O+ + OH- H2O

Ex: 2NaOH + H2SO4 Na2SO4 + 2H2O or HCl + NaOH NaCl + H2O9. In neutralization reactions for every mole of H+ that reacts, one mole of OH-

reacts. Ex: If 2 moles of H+ ions are neutralized, how many moles of OH- ions are needed? Ans. = 2 moles

10. Formula for titration of acids and bases : when different concentrations of an acid and base are mixed, we can use the titration formula to find an unknown (a version of the formula is also found on Table T, just add #H+ and #OH- to it):

(#H+) MA VA = MB VB (#OH-)MA = molarity of the acid; VA = volume of the acid; #H = number of H+ in acid formulaMB = molarity of the base; VB = volume of the base; #OH = number of OH- in base formula

Example: What volume of 2.0M NaOH is needed to neutralize 30. mL of 4.0M H2SO4?Answer: (2)(4.0M)(30.mL) = (2.0M)(VB)(1) VB = 120mL11. Definitions of Acids and Bases:

a. Arrhenius Theory: acids yield H+ ions, bases yield OH- ions. Stronger acids produce more H+ ions than weaker acids. Strong bases produce more OH- ions than weak bases.

b. Bronsted-Lowry Theory: acid = proton (H+ donor); base = proton acceptor. Ex. NH4+ is a BL acid: NH4+ + H2O NH3 + H+ (NH4+ donates proton to H2O)Ex. NH3 is a BL base: NH3 + H2O NH4+ + OH- (NH3 accepts proton from H2O)Amphoteric – a substance that can either accept or donate H+ depending on the conditions. Ex: H2O can become either H3O+ or OH-

11. The larger the [H+], the stronger the acid, the lower the pH, and the better the conductivity.

12. Strongest bases have Group 1 with OH Ex – NaOH, KOH, etc.13. Acids react with any metal higher than H2 on Reference Table J spontaneously.

Ex: Cu will not react with an acid, but Mg will. Metal + Acid Salt + Hydrogen gas

Mg + HCl MgCl2 + H2

29


30


Topic #12 – Oxidation-Reduction 1. Know rules for determining oxidation #’s (oxidation numbers are found on

periodic table. 2. The sum of the oxidation numbers in a compound is zero. Elements by

themselves have an oxidation state of zero. Ex: Na0, Mg0, Cl20. Ex: Find the oxidation number of S in Na2S2O3. Ans: Na2+1S2+4O3-2.

3. Oxidation= a loss of electrons. Reduction= a gain of electrons. Remember: LEO the lion goes GER = “Losing Electrons is Oxidation” and “Gaining Electrons is Reduction” andOIL RIG = “Oxidation is Losing electrons” and “Reduction is Gaining electrons”.

4. Oxidation Number:a. Oxidation = increasing oxidation numberb. Reduction = decrease in oxidation number.Ex: Mg0 + Zn0Cl2+2 Mg+2Cl2-1 + Zn0. Mg0 (0+2) is oxidized and the reducing agentZn+2 (+20) is reduced and the oxidizing agent.

5. In a redox reaction, there must be a change in oxidation numbers. In a double replacement there is no redox, there is always a redox reaction in a single replacement reaction (change in oxidation numbers).

6. As the number of particles oxidized during a reaction increases, the number of particles reduced also increases.

7. Know the reaction: Cu + 2 AgNO3 Cu(NO3)2 + 2 AgNet Reaction: Cu0 + 2 Ag+1 Cu+2 (blue color) + 2 Ag0

NO3- is the spectator ion in the above reaction (oxidation state did not change).

8. Writing half reactions:a. Equations must be balanced according to both mass and charge.

9. The number of each element must be the same on both sides of the arrow, and the sum of the charges must be same on both sides of the arrow. Ex:

31


10. Some substances (Ex: Sn+2) can act as both an oxidizing and reducing agent:a. Sn+2Sn+4 + 2 e- Sn+2 is oxidized = reducing agentsb. Sn+2 + 2e- Sn0 Sn+2 is reduced = oxidizing agent

11. To find the number of moles of electrons transferred, find the number of moles of e- transferred from oxidized to reduced part of equation.

Ex: Find the total number of moles of e- transferred from 2Al(s) to 3 Cu+2 in:2 Al(s) = 3 Cu+2(aq) 2 Al+3 + 3 Cu(s)

Use Half Reaction:Al0Al+3 + 3 e- therefore 2 Al0 2Al+3 + 6e- which means 6 moles of e-

transferred 12. Reference Table J shows Metals from most to least

likely oxidized and lose electrons; and nonmetals from most to least likely reduced and gain electrons.

13. Electrochemical Cell (Uses chemical redox reaction to create electricity):

32


a. Produces electricity by means of a spontaneous reactionb. Be able to write half-reactions for the cell. The more active metal (Closer

to top of Table J) is oxidized, the less active metal (closer to bottom of Table J) is reduced.

c. Electrons flow from the more active metal to the less active metal.d. Oxidation occurs at the negative electrode (anode)e. Be able to write overall reaction for the cell and the moles of electrons

transferredf. The salt bridge functions to allow the migration of ions. If the salt bridge

is removed the cell voltage = 0 volts.“An Ox, Red Cat” = Anode is oxidized, cathode is reduced.”

14. Electrolytic Cell (Electricity added produces redox reaction)

a. Uses electricity to force a non-spontaneous reaction to occur

b. Oxidation occurs at the positive electrode (anode)

c. Be able to write half reactions at each electroded. In electrolysis a compound is broken down into

its free elements. Ex: Hoffman Apparatus- 2H2O 2H2 + O2.

“Jump start a car – needs a battery, Nephew, have you ever seen a Red Cat? NEGATIVE.Have you ever seen an ox? POSITIVEly.”

For Both types of cell: “An Ox, Red Cat” – anode is oxidized, cathode is reduced.

15. Predicting if a redox reaction will occur:a. Use Table J: If the free element is higher than the first element in the

compound, the reaction will take place. If not it will not take place.Ex: 2 Al +3 CuCl2 3 Cu + 2AlCl3; the reaction will take place since Al is higher up than Cu.

Sayings to Remember for Redox Reactions: “LEO the lion goes GER” – loss of electrons is oxidation “OIL RIG” - oxidation is losing electrons, reduction is gaining electrons “An Ox and Red Cat” – anode is oxidized, cathode is reduced.

33

http://www.bing.com/images/search?q=leo+lion+ger&view=detail&id=F0B2CD3C7A68A8C44D6500B4262E83A4825AE2FC&first=0&qpvt=leo+lion+ger&FORM=IDFRIR


Topic #13 – Organic Chemistry Anything organic has to have the element CARBON. All carbon atoms need four bonds. Hydrogen and Halides (group 17)

may combine with Carbon and can form only 1 bond. Properties of Organic Molecules:

a. Non-electrolytes (except acids “COOH”).b. Low boiling and melting points (Ex – Sugar C6H12O6)c. Compounds usually insoluble in waterd. Compounds react slowlye. Are molecular in structure

Homologous Series – successive members differ by CH2 group Hydrocarbon – contains both Carbon and Hydrogen only. Ex – hexane;

propene; 2 heptyne. Saturated = all single bonds; Unsaturated = at least one double or

triple bond. Table P gives organic prefixes, Table Q summarizes alkanes, alkenes,

alkynes (“n” = # of carbons). Alkanes : CnH2n+2 - contain only single bonds (saturated). Ex: CH4, C2H6,

C3H8, C4H10, etc. Names end in “-ane” = methane, ethane, propane, butane…

Alkenes: CnH2n – contain one double bond (unsaturated); starts with ethene (no meth-). Ex:C2H4, C3H6, C4H8, C5H10, etc. Names end in “ene” = ethene, propene, butene, pentene…

Alkyne: CnH2n-2 – contain one triple bond (unsaturated). Starts with ethyne (no meth-). Ex: C2H2, C3H4, C4H6, etc. Name ends in “-yne” = ethyne, propyne, butyne…

As the molecular mass of organic compounds in each series increases, the boiling point increases due to Van der Waals forces.

Alkyl Radicals : formed from each alkane by removing H. Ex: methyl group CH3

Ethyl group - C2H5 . Naming Organic Compounds:

a. find longest consecutive carbon chainb. find closest side chain to first carbonc. use prefix for more than one of same side chain with di-, tri-, tetra-

etc.d. End with suffix of parent. Ex –ane for alkane, -ene, for alkene, -yne

for alkyne.

Example: 3,4 dimethyl heptane and 2,2 diethyl-4-methylpentane Isomers: compounds that have the same molecular formulas but

different structural formulas (different arrangement of atoms.) For hydrocarbon molecules you need at least four carbon atoms to have an

propene

propane

ethyne

34


isomer. The more carbons, the more possible isomers (Isomers have different properties). Ex –

15. Functional Groups: (Use Table R). Know how to draw out each example in Table R:

Alcohols: R-OH. Monohydroxyl (1-OH), dihydroxyl (2-OH), trihydroxyl (3-OH)

a. Primary Alcohol: the “C” attached to the “OH” has 1 Carbon directly attached to it (on the last “C”)

b. Secondary Alcohol: the “C” attached to the “OH” has 2 “C’s” directly attached to it.

c. Tertiary Alcohol: the “C” attached to the “OH” has 3 “C’s” directly attached to it

Some Functional Group Sayings to help you remember them: Organic Acids: they’re “COOH”. Ketone: Think of a dog howling “O” straight up. (Means

“O” in middle has a double bond.

Ester: “ ” is she ugly, “O” is she ugly (to side means single bond, up down is double bond.” She is a COOC

Ether: or there is always a COC Aldehyde: “Al the cheerleader saying the H and O are

“hyding” at the end. Al dates a lot of women at the end –HO”.

Amine: Nitrogen but No Oxygen. Amide: Nitrogen Does have Oxygen16. Organic Reactions Addition: involves unsaturated hydrocarbonsEx: C2H4 + Cl2 C2H4Cl2. Alkenes are unsaturated, and just like in math, one answer when you add. Substitutions: involves saturated hydrocarbons. (S in Saturated, S in

substitution).Ex: C2H6 + Cl2 C2H5Cl + HCl. Alkanes are saturated, and there are two products. Like sports, one thing is switched with another. Esterification: acid + alcohol ester + H2O. Saponification: can’t spell saponification w/o SOAP. Fat + Base soap

+ glycerol.

Pentane

Boiling Point = 36°C

Methyl Butane


2,2 Dimethyl Propane


35


Fermentation: C6H12O6 2C2H5OH + CO2. Polymerization: Small chains put together to form long chains. This

forms natural compounds such as nylon, rayon, starches, cellulose; and artificial compounds such as Teflon, polystyrene, etc.

a. Addition Polymerization: joining monomers (small chains) of unsaturated compounds by opening n(C2H4) (C2H4)n. Chains linked together.

b. Condensation Polymerization: long chains formed and water as a byproduct. “Dehydration synthesis”.

36

· Web viewChemistry Regents Review Concept Sheets Table of Contents Topic 1 – Math...

Documents

Transcript of · Web viewChemistry Regents Review Concept Sheets Table of Contents Topic 1 – Math...