CHAPTER 10 MODERN ATOMIC THEORY

• Atoms emit colors when electron enter the excited state and leap to a higher energy level.

• As they return to the ground state, they release that energy a a photon or light particle.

• Electrons can act as a particle or a wave. • This behavior led to the modern atomic

theory of wave mechanics.• So, atomic structure went through four main

stages to this point:– Plum pudding– Rutherford’s nuclear model– Bohr orbit model– Wave mechanic model

Plum pudding

• Lord Kelvin’s model• Protons are scattered in a negative “pudding”

of electrons

Nuclear Model

• Rutherford’s model• Moves all protons to a nucleus• Electrons are scattered in a cloud around the

nucleus

Bohr Orbit

• Bohr’s model• Keeps protons in a nucleus• Moves electrons into orbits around the

positive charge• Works only for hydrogen

Wave Mechanic Model

• Also known as the orbital model• Keeps protons in the nucleus• Moves electrons in specific areas known as

orbitals around the nucleus• Orbitals are different shapes• Electrons behave in orbitals according to the

Pauli Exclusion Principle

Pauli Exclusion Principle

• Only two electrons per orbital• These electrons must spin in opposite

directions• Electrons must add one at a time before

pairing up

Electron configurations

• Electrons are ordered in wave mechanic model– Energy levels– Orbital– Orbital orientations allow for total electron count

• Energy levels correspond to the row element is in the periodic table

• Orbitals make a pattern in the table• Each orbital can be in different 3D spots, so you

can have more than one orbital per energy level

• s orbital can hold 2 electrons– Sphere with one orientation

• p orbital can hold a total of 6 electrons– Bow tie with three orientations

• d orbital can hold a total of 10 electrons– Complex with 5 orientations

• f orbital can hold a total of 14 electrons– Complex with 7 orientations

Atomic electron configurations

• Superscripts add up to element number• Must follow arrangement of orbitals on the

periodic table• Ions have ec as well– Strive to be like noble gases– Either gain electrons to fill shell or lose electrons

to fall back one energy level

• Hydrogen– 1s1

• Neon– 1s22s23p6

• Write electron configurations for – Ca (20)– I (53)– F-1 (9)– Mg+2 (12)

Valence electrons

• Electrons in outermost orbital or shell• Only electrons used in bonding and reactions• Electrons in highest energy level in electron

configurations• Determine physical characteristics of atoms

and lead to trends in the periodic table

Trends in the periodic table• s,p,d,f orbital arrangement• Metal, non-metal, metalliods• Atomic size– Increase as you move down table– Decrease as you move across table

• Ionization energy– Energy needed to remove electrons– Decrease as you move down table– Increase as you move across table

• Electronegativity– Desire for an element to pull electrons toward itself– Decrease as you move down table– Increase as you move across table

CHAPTER 11 BONDING

Electronegativity Values

• All elements have an electronegativity value• The difference in these values determines the type

of intramolecular bond (within the molecule)– Non-polar (0-0.7)– Polar (0.7-1.5)• 1.5-2.0 depends on the element as metal or non-metal• Metal/non-metal are ionic• Non-metal/non-metal are polar

– Ionic (> 2.0)• All metal/non-metal bonds are ionic no matter what Δen

Intramolecular Bonding • Non-polar– Equal sharing of electrons

• Polar – Unequal sharing of electrons– Creates dipoles

• Ionic – No sharing

• Intramolecular bonding influences intermolecular bonding (between molecules)

Intermolecular bonding• London dispersion– Non-polar

• Dipole-dipole– Polar

• Hydrogen bonding– Polar bonding between molecules that have H and O, N,

and F• Ionic– Large networks created– Alternating positive and negative charges

• Metallic– Sea of electrons– Conducts heat and electricity

Representing bonding

• Lewis structures are visual ways to represent bonds– Use only valence electrons– Shared and unshared pairs of electrons– Dashes represent one shared pair of electrons

Molecular geometry• Bonding arrangement leads to the geometry

of the molecules• Lewis structures can help predict molecular

geometry• Three electron pair arrangements– Linear– Triangular– Tetrahedral

• Use the chart to see molecular geometry

Total number of electron pairs

Electron pair arrangement

Shared or bonded pairs of electrons

Lone pairs of electrons

Molecular geometry

2 Linear 2 0 linear

3Triangular 3 0 Triangular

2 1 V-shape or bent

4

Tetrahedral 4 0 Tetrahedral

3 1 Trigonal pyramid

2 2 V-shape or bent

CHAPTER 7 REACTIONS

Classes of Reactions

• Precipitate– A reaction where a solid forms from the mixing of two

clear liquids• Oxidation reduction– The movement of electrons

• Oxidation is loss reduction is gain (OIL RIG)– Combustion reactions

• Hydrocarbon reacts with O2 to make CO2 and H2O

• Acid base reactions– HA plus BOH react to make a salt and water

Types of reactions• Synthesis – A + B → AB

• Decomposition– AB → A + B

• Single displacement– AB + C → AC + B

• Double displacement– AB + CD → AD + CB

Predicting reaction products• Using the general equations on the previous slide,

you must be able to predict the products of a reaction

• Remember to make neutral compounds first.• For example:– C + O2 →– C + O2 → CO2

– HCl + NaOH →– HCl + NaOH → H2O + NaCl– K + CuSO4 →– K + CuSO4 → K2SO4 + Cu

Precipitate reactions

• Create a solid• Use the flow chart to identify products that

should be solids

SO4-2 salts

Aqueous EXCEPT

with Ba+2, Pb+2,

Ca+2

Hydroxide (OH-1) salts

Solid EXCEPT

with K+1, Na+1, Ba+2, Ca+2

K+1, Na+1, NH4

+1, NO3-1

Always aqueous (no matter what

partner)

S-2, CO3-2, PO4

-3

Solid EXCEPT

with K+1, Na+1,

NH4+1

Halide salts (Cl-1, Br-1, I-1)

Aqueous EXCEPT

with Ag+1, Hg2

+2, Pb+2

NaS (aq) + CaNO3 (aq) →

Na+1 (aq) + S-2 (aq) + Ca+2 (aq) + NO3-1 →

NaS (aq) + CaNO3 (aq) → CaS (s) + NaNO3 (aq)

NET: S-2 (aq) + Ca+2 (aq) → CaS (s)

Oxidation Reduction

• Electron transfer

2 Fe + 3 CuSO4 → Fe2(SO4)3+ 3 Cu• Fe begins at Fe0 and ends up as Fe+3

– Lost electrons• Cu begins as Cu+2 and ends up as Cu0

– Gains electrons

Half Reactions

2 Fe + 3 CuSO4 → Fe2(SO4)3+ 3 Cu Fe0 → Fe+3 + 3e-

Cu+2 + 2 e- → Cu0

2Fe0 →2 Fe+3 + 6e-

3Cu+2 + 6e- →3 Cu0

2Fe0 + 3 Cu+2 → 2Fe+3 + 3 Cu0

Solution reactions

• Dissociation– Process where an ionic compound separates in

water into ions– Dissociated ions are known as electrolytes• Electrolyte solutions can conduct electricity

– Acids dissociate into H+1 ions and anions• HF → H+1 + F-1

– Bases dissociate into cations and OH-1

• NaOH → Na+1 + OH -1

CHAPTER 12 GASES

Kinetic molecular theory of gases• Gases are made of tiny particles• Particles are very small and far apart so their

size is zero• Particles are in constant, random motion and

collide to create pressure• Particles do not repel or attract each other• Kinetic energy of the particle is directly

related to the temperature

Gas Laws

Law Variables Constants Proportional

Avogadro’s V, n T, P Directly

Boyles P, V T, n Indirectly

Charles V, T P, n Directly

Gay-Lussac P, T V, n Directly

Combined P, V, T n N/A

Ideal n, V, P, T R N/A

• Dalton’s law of partial pressure– Pt = P1 + P2 + P3…….– Key is to find the pressure with PV=nRT– Total pressure is the sum of the individual

pressures

Gas Stoichiometry• All gas laws and gas stoichiometry use Kelvin

scale for temperature• 1 mole of any gas (the volume) is equal to 22.4 L• STP is standard temperature and pressure– 1 atm and 273 K

• Problems must always focus on finding moles first, then volume– If at STP, use 22.4 L– If not at STP, use PV=nRT once you have moles (n)