Post on 29-Dec-2015
F. BASIC STRUCTURES OF MATTER
1. ATOMS– basic building blocks of matter– STRUCTURE:• protons (p): positively charged
sub-atomic particle• electron (e-): negatively charged
sub-atomic particle• neutron (n): electrically neutral
F. BASIC STRUCTURES OF MATTER
A ZM CMass Number
Atomic Number
Charge
Atomic Symbol
M = (# of p+) + (# of n0)C = (# of p+) - (# of e-)A = # of p+
Atoms of the same element having the same number of protons (same atomic number) but
different number of neutrons and different atomic mass are called isotopes.
ORBITALS AND QUANTUM NUMBERS
• Quantum numbers: describe certain aspects of the atom
• Orbital: specific distribution of space orientations between electrons
• Electron shell: orbitals with the same value n/1st or principal quantum number
• Subshell: same set of n and l (2nd quantum number)
TYPES OF QUANTUM NUMBERS1. First/Principal (n): energy level of the electron (the
higher, the more energy/size)[for the maximum number of orbitals = n2)
2. Secondary/Azithumal (l = 0 until n-1): sublevel within the energy level = shape of the of the orbital
[s = 0 = spherical, p = 1 = dumbell, d = 2 = clover , f = 3 = more complex]
3. Third/Magnetic (ml): the particular orbital and its space orientation [ s = 1, p = 3, d = 7, f = 10]
4. Fourth/Spin (ms): direction of the spin [+ ½ or – ½ ]
RELATIONSHIP BETWEEN QUANTUM NUMBERS
n l Subshell designation Ml
Number of Orbitals in a subshell (n2)
Total number of electrons
1 0 1s 0 1 2
20 2s 0 1 2
1 2p -1, 0, 1 3 6
3
0 3s 0 1 2
1 3p -1, 0, 1 3 6
2 3d -2, -1, 0, 1, 2 5 10
4
0 4s 0 1 2
1 4p -1, 0, 1 3 6
2 4d -2, -1, 0, 1, 2 5 10
3 4f -3, -2, -1, 0, 1, 2, 3 7 14
PAULI’S EXCLUSION PRINCIPLE
• A maximum of two electrons can occupy an orbital (with different spins)
• No two electrons can have the same quantum numbers
ELECTRON CONFIGURATION
PERIODIC LAW
• Physical and chemical properties of elements are the periodic functions of their atomic number
• Elements with similar properties = similar arrangement of outer shell electrons/same group
• Valence electrons: found in the outermost shell of elements in their ground state
PERIODIC TRENDSINCREASES GOING DOWN/DECREASES LEFT TO RIGHT
DECREASES GOING DOWN/INCREASES LEFT TO RIGHT
Atomic/Ionic Radius/SizeElectronegativity
(Ability of an element to attract other electrons to itself in a covalent bond)
Metallic Properties (what makes them metallic – shiny, malleable, etc.)
Ionization Energy(Energy required to remove an
electron from the outermost shell)Electron Affinity
(energy released when an electron is added)
• Molecules: two or more atoms tightly bonded together• Ion: atom with a gained or loss electron (anion: negatively
charged; cation: positively charged)
INTRAMOLECULAR FORCES OF ATTRACTION
1. Covalent bond: between 2 atoms; by electron sharing (nonmetallic + nonmetallic)
2. Ionic bond: between + and – ions; form crystal structures; by electron transfer (nonmetallic + metallic)
3. Metallic bond: between free-floating valence electrons = high conductivity and luster (metallic + metallic)
VAN DER WAAL’S FORCES OF ATTRACTION (IMFA’s)
• Between atoms of different molecules1. London dispersion forces: weakest, between nonpolar
atoms and molecules (nonpolar + nonpolar) (CH3-CH3)
2. Dipole-dipole forces: between polar molecules (CO)3. Hydrogen bond: between a hydrogen atom of a polar
molecule + (any) F/O/N atoms in another polar molecule (H20, NH3)
4. Ion-dipole forces: strongest, between polar molecules/ions in solutions (ionic molecules + ionic salts/polar solvents)
WRITING CHEMICAL FORMULAS
• Chemical formulas: denote the number and kinds of atoms in a compound (reacting substances)
1. Represent with symbols (cation first, then anion).2. Indicate oxidation states (to the upper right, put
parentheses for polyatomic ions).3. Write the subscript equal to the oxidation number
of the other element (‘switch’) [omit the ‘1’s’].4. Reduce subscripts to their lowest terms.
FORMULAS TO REMEMBER!
1. FORMULA WEIGHT/MASS or MOLECULAR WEIGHT/MASS
– sum of all atomic weight (AW) (g/mol)
2. PERCENTAGE COMPOSITION
FORMULAS TO REMEMBER!
3. EMPIRICAL AND MOLECULAR FORMULAa.) convert mass in grams to mole ** assume 100 g of any sample so that % composition = value of mass in gb.) get the smallest whole number ratio of the
elements ** divide mole value of each element by the smallest mole value
THIS WILL BE YOUR EMPIRICAL FORMULA (EF)
c.) ** get EFM the same way you get FM
FORMULAS TO REMEMBER!
4. STOICHIOMETRY ** Avogrado’s number = 6.02 x 1023 molecules ** Molar Mass = mass in grams of 1 mole of a substance
CONVERSION
GRAMS MOLES MOLECULES
FORMULAS TO REMEMBER!
4. REACTION STOICHIOMETRY CONVERSION use molar mass of A use molar mass of B
grams of moles of moles of grams ofsubstance A substance A substance B substance B
use coefficients of A and B from the
BALANCED EQUATION
FORMULAS TO REMEMBER!6. GAS LAWS
a.) Boyle’s Law
b.) Charles’ Law
c.) Gay- Lussac’s Law
d.) Combined Gas Law
e.) Avogrado’s Law
f.) Daton’s Law … +
g.) Graham’s Law
h.) Ideal Gas Law
R = 0.0821 L-atm/mol-K
FORMULAS TO REMEMBER!
7. GAS DENSITIES AND MOLAR MASS
At STP,
standard temperature: 0°C = 273.15 K = 32°Fstandard pressure: 1 atm = 760 torr = 760 mmHgmolar volume of gas: 1 mol = 22.4 L
This states that ideal gases: 1. the movement of particles is in continuous
random motion 2. pressure of gas in due to the bombardment of
the molecules to the container 3. collisions between or among particles are
elastic 4. the kinetic energy of the system is
proportional to temperature 5. the wide separation between molecules cause
the attractive and repulsive forces to be negligible
KINETIC MOLECULAR THEORY OF GASES
1. A substance composed of two or more elements chemically united is called
a) an isotopeb) an elementc) a mixtured) a compound
2. When energy, like light and heat, is liberated during a reaction such as burning of fuels, this type of reaction occurs.
a) endothermicb) nuclear reactionc) exothermicd) both a and b
3. Chemical action may involve all of the following EXCEPT
a) combination of atoms of all elements to form a moleculeb) breaking down of compounds into elementsc) reacting a compound and an elementd) separation of the components of a mixture
4. When decomposed chemically, 73 g of a sample of HCl produces 71 g of Cl2 and 2 g of H2, while 34 g of H2S sample produces 32 g of S and 2 g of H2. This is an example of:
e) Law of Conservation of Massf) Law of Thermodynamicsg) Law of Multiple Proportionsh) Law of Definite Proportions
5. The theory states that no 2 electrons within the same atom can have the same set of quantum numbers.
a) Hund’s Ruleb) Aufbau’s Principlec) Pauli’s Exclusion Theoryd) Dalton’s Rule
6. Elements with the most stable configuration belong to:
e) transition metalsf) inert gasesg) alkali metalsh) halogen family
7. In the periodic table, which can also indicate the highest energy level or the highest principal quantum number of an atom?
a) groupb) atomic numberc) periodd) atomic mass
8. The total number of orbitals in the fourth energy level is:
e) 4f) 16g) 8h) 18
9. Sodium chloride, most commonly known as table salt, has a pH of:
a) = 7b) < 7c) > 7d) = 0
10. What do you call the solid left after the separation of mixtures?
e) solventf) precipitateg) mother liquorh) supernatant
11. From the given statements below, which is INCORRECT?
a) All atoms of a given element are identical.b) Atoms combine in small whole number ratio.c) All atoms change their chemical identity during a chemical
reaction.d) Elements can combine to form compounds.
12. Isotopes, which are the reasons why elements have no whole mass numbers, have:
e) the same number of protons but different number of electrons
f) the same atomic number but different mass numbersg) the same number of protons but different number of
neutronsh) both b and c
13. The intermolecular forces that exist between the molecules of CCl4, a non-polar compound, is/are:
a) dipole-dipoleb) London forcesc) Hydrogen bondd) both a and b
14. If Magnesium forms a compound with chlorine, what is the formula of the molecule?
e) Mg2Cl
f) MgCl2
g) Mg2Cl3
h) MgCl3
15. Which of the following statements is FALSE for a neutral atom?
a) number of neutrons is not always equal to the number of protons
b) number of electrons in the atom is equal to the number of protons
c) nucleus has a positive charged) nucleus has a neutral charge
16. The maximum number of electrons that can be placed in the second principal energy level, n=2, of an atom is:
e) 6f) 8g) 4h) 2
17. Why are you advised not to heat a tightly closed vessel?
a) heat increases the volume of the vesselb) pressure might decrease causing the vessel to
implodec) pressure will build up causing the vessel to exploded) heat might be trapped inside the vessel
18. Gases will approach ideality (ideal state) at:e) low pressure and high temperaturef) low pressure and low temperatureg) high pressure and high temperatureh) high pressure and high temperature
19. What should be the last 2 coefficients when the following reaction is balanced?
FeS2 + O2 ___ Fe2O3 + ___ SO2
a) 2, 3b) 2, 4c) 2, 6d) 2, 8
20. Gas flows from an area of _____ pressure to an are a _____ pressure.
e) higher-lowerf) equal-equalg) medium-highh) lower-higher
21. What type of reaction occurred in the following:4 NiCO3 + O2 2 Ni2O3 + 4 CO2 + heat?
a) double displacementb) single replacementc) combustiond) decomposition
22. The reaction of ethylene, C2H4(g) and hydrogen chloride, HCl(g) to form ethyl chloride C2H5Cl(g) is an example of what type of reaction?
e) substitutionf) synthesisg) neutralizationh) decomposition
23. A real (or non-ideal) gas will:a) have a higher kinetic energy than ideal gasesb) have a higher boiling than ideal gasesc) not have negligible attractive and repulsive forces
between its moleculesd) conform to the kinetic molecular theory of gases
24. The sum of the partial pressures of gas components is equal to the total pressure of the system. This is stated in which law?
e) Boyle’s Lawf) Combined Gas Lawg) Dalton’s Lawh) Gay-Lussac’s Law
25. Which of the following are not STP conditions for 1 mole of ideal gas?
a) 760 torrb) 273 °Cc) 0°Cd) 22.4 L
26. Which of the following is not stated in the Kinetic Molecular Theory of Gases?
e) temperature of gas is related to the kinetic energy of the molecules
f) pressure is due to the bombardment of molecules to walls of container
g) the larger the molecule of a gas, the greater its kinetic energy
h) collisions are elastic
27. Which has more atoms: 100 mol of Cu, 100 mol of Ag, 100 mol of Au, or 100 mol of Hg?
a) Hgb) Auc) Cud) none of the above
28. How many moles of oxygen atoms are there in 3 moles of Ca3(PO4)2?
e) 8f) 4g) 24h) 36
EXERCISES!ATOM # of n0 # of p+ # of e- C M
72 51
-2 127
n0 = M - A = 127 – 52 = 75
p+ = A =52
e- = p+ – C = 52 – (-2) = 54
EXERCISES!ATOM # of n0 # of p+ # of e- C M
72 51
-2 127
4 197
75 52 54
n0 = M - A = 197 – 79 = 118
p+ = A =79
e- = p+ – C = 79 – 4 = 75
EXERCISES!ATOM # of n0 # of p+ # of e- C M
72 51
75 52 54 -2 127
118 79 75 4 197
3
M = n0 + A = 72 + 51 = 123
e- = p+ – C = 51 - 3 = 48
EXERCISES!Some AW in amu or g/mol
H = 1.0 Na = 23.0 N = 14.0 C = 12.0 O = 16.0 Al = 27.0
• Find the formula weight of:– H2O
= 2 (AWH) + AWO
= 2 (1) + 16 = 18 g/mol– NaOH
= AWNa + AWO + AWH
= 23 + 16 + 1 = 40 g/mol– Al(NO3)3
= AWAl + 3 (AWN) + 9 (AWO)
= 27 + 3 (14) + 9 (16) = 213 g/mol
EF: CH3
EXERCISES!given: MWcompound = 30 g/mol
80% C ; 20% H (assume 100g)
divide by their AW
• MF = MF = MF = 2
MF: 2 (CH3)
MF: C2H6
ATOMS MASS MOLE RATIOCH
80 g20 g
( 6.67( 20
( 1( 3
by smallest mole value
EXERCISES!Some AW in amu or g/mol
H = 1.0 Na = 23.0 O = 16.0
• 1.5 mol NaOH = ____ g NaOH = ____ molecules NaOH
** FW = 23 + 16 + 1 = 40 g/mol– 1.5 mol NaOH
= 60 g NaOH– 1.5 mol NaOH
= 9 x 1023 molecules NaOH