Chapter 4
Subatomic particles: Protons (+), neutrons (0), electrons (-)
Atomic mass = Protons + neutrons
Ionic charge = protons – electrons
Atomic number = # of protons
Types of radiation: alpha (4/2 He) stopped by paper/cloth
Beta (0/-1 e) stopped by sheet of metal
Gamma (0/0) stopped by 1-foot of concrete
Dalton’s Atomic Theory:
All atoms of the same element are the same (isotopes disprove)
Atoms cannot be created or destroyed (nuclear radiation disproves)
Molecules rearrange in whole number ratios in a chemical reaction
All matter is made of atoms
Atoms combine in whole numbers to form compounds
Rutherford gold foil experiment (proved the existence of a positive nucleus)
Chapter 5
Bohr – ring model of the atom (quantum energy levels for electrons)
Quantum – amount of energy gained or lost by an electron in an atom (a discrete packet of energy)
Photon – a packet of light emitted from an electron losing energy
Photoelectric effect – discovered by Einstein (Nobel Prize!) – describes how energizing an element causes it to light up.
Atomic emission spectrum – different elements present different wavelengths of light due to different arrangement of electrons
deBroglie – wave equation
Schrodinger – wave-particle duality
Heisenberg – uncertainty principle
Quantum-mechanical model of the atom – electrons not only exist on specific energy levels but also in vague clouds of probability
Electron Configurations:
1s^2, 2s^2, 2p^6…etc… [electron rain]]lji-
Orbital diagrams (that thing with the boxes and arrows)
Aufbau Principle
Pauli’s Exclusion Principle
Hund’s Rule
Chapter 6- Periodic Table
Period – energy level
Group – valence electrons
Valence electrons – outermost electrons (partake in reaction)
s-, p-, d-, f- blocks
Mendeleev – organized modern periodic table by mass
Newlands – Octet rule
Moseley – corrected periodic table by organizing by atomic number
Representative elements – s, p blocks
Transition metals – d,f blocks
Inner transition metals – f block
Alkali metals – Group 1
Alkaline earth metals – Group 2
Halogens – Group 7
Noble gases – Group 8
Periodic Table Trends:
Atomic radius – increases to bottom left corner
Electronegativity – increases to top right corner
Ionization energy – increases to top right corner
Ionic Radius – gaining electrons increases size, losing electrons decreases size
Chapter 8 – Ionic Compounds
Ions result from gain/loss of electrons
Cation – positive ion
Anion – negative ion
Opposite charges attract, like charges repel
Ionic Bond results from electrostatic attraction between ions.
Lattice energy – ionic bond strength, related to electronegativity trend
Crystal Lattice – structure of ionic solids
Electrolyte – ionic compound dissolved in water is a conductive substance
Metallic bonds
Sea of electrons/ delocalized electrons
Alloy- mix of metals with other elements to create new physical
properties
Differences between ionic v. metallic solids
Ionic:
Hard, brittle, high melting point
Metallic solids:
Soft, flexible (malleable, ductile), low melting points, conductive
Differences due to differing electron arrangements: ionic is highly rigid structure versus metallic is very loose.
Chapter 9 – Covalent Bonds
Covalent bond: when electrons are shared between two atoms
Lewis structures: extension of electron dot structure. (the lines and dots)
1) ID number of valence electrons/ number of bonds that can be formed
2) Most bonded atoms go in middle, less bonded on outside
3) Draw bonds from outside and work inwards
4) Dots = Group # – bonds formed
Naming covalent molecules:
Use numerical prefixes (mono-, di-, tri-, etc…
Less electronegative element named first
No use of “mono-“ for first element named (Ex: carbon dioxide, NOT monocarbon dioxide)
Sigma bonds are the first/single bond between two elements
Pi bonds: the additional bonds after sigma bond (Ex: a triple bond has one sigma & 2 pi bonds)
Strength of covalent bonds depends on how close the elements are to each other. The closer they are, the stronger the bond. Double & triple bonds are closer/stronger than single bonds.
Single< double< triple (Bond Strength)
Naming Acids:
Binary Acids (H-X): hydro-element-ic acid
Ex. Hydrochloric acid
Oxyacids (HXO): ends in –ite then –ous acid (H2SO3: sulfurous acid)
-ate then –ic acid (H2SO4: sulfuric acid)
VSEPR Models: pg. 260
Number of lone pair electrons + bonded elements counts up to VSEPR hybrid type.
Ex: NH3 : sp3
Polarity: results from one element attracting electrons more strongly in a molecule than others.
More electronegative elements attract electrons more strongly.
Ex. H2O: Oxygen is more electronegative than hydrogen so the oxygen side has more electrons and is slightly more negative than the hydrogens.
Properties of Covalent Compounds:
Can exist in any state at room temperature, depends on structure of the molecule.
Covalent Network Solid: strongest type of covalent bonding: Where atoms are bonded atop, below, and to the sides. Makes strongest substances on the planet. Ex. Diamond v. graphite/pencils
Chapter 10 – Chemical Reactions
All chem. Reactions are comprised of 2 parts, reactants and products
Reactant: stuff you start out with (before reaction)
Product: stuff you finish with (after reaction)
Skeleton equation: basic layout of compounds in the reaction
Balanced equation: number ratios of compounds stated (coefficients)
*Pro-tip* balance oxygen last
In combustion reactions water needs to have an even coefficient.
Types of Chem Reactions:
Synthesis : A + B à AB
Decomposition: AB à A + B
Single Replacement: AX + B à A + BX
Double Replacement: AX + BY à AY + BX
Combustion: A + O2 à AO + energy
Precipitate: solid produced during a chemical reaction.
Phase labels: (s) solid
(l) liquid
(g) gas
(aq) aqueous: dissolved in water
Ionic equations: (complete & net)
1) ID the precipitate formed
2) Break apart the compounds into their ions
3) ID the spectator ions (ions that do NOT produce precipitate)
4) Cross out spectator ions
5) Rewrite equation without spectator ions
Chapter 11
Mole: number of atoms found in 12 grams of Carbon
Avogadro’s number: 6.02 x 10^23 particles
Molar mass: mass amount for 1 mole of a compound.
Ex. Water = H2O = 1+1+16= 18 grams/mole
Converting between particles à molesà grams
1 mole = 6.02x10^23 particles
1 mole = molar mass in grams
· Going towards moles then divide
· Going towards molar mass or # of particle then you multiply
Percent Composition:
__Mass of element______x 100 = % comp.
Total mass of compound
Do for each element in compound
Empirical Formula:
1) divide each element from atomic mass
2) compare which has lowest result
3) divide all results with the lowest result
4) write down the ratio
Molecular formula:
1)Molar mass given = n
Empirical formula mass
2) multiply ratio numbers by n
hydrate: compound that binds to water
Chapter 12 - Stoichiometry
Understand how to convert particles or grams to moles
Going from Moles A to Moles B uses the balanced equation (coefficients used)
Limiting Reactant: reactant amount that produces less product (remember hotdogs v. buns)
% yield: ____actual____ x 100 = % yield
Theoretical
Chapter 13.1 & - 14 Gas
3 states of matter differ by energy amount
Least - Solids< liquids< gas – most energy
Kinetic Molecular Theory:
Particles are small relative to distances between particles
Particles do not stick to each other (elastic collisions)
Particles experience an average velocity
Temperature: average kinetic energy of particles
Pressure: Results from particles bouncing off walls of container
Units: Pressure : atm, mmHg, torr, kPa
Volume: Liters
Temperature : Celsius, Kelvins
K = C + 273
STP: (Standard Temperature Pressure) : 0 Celsius or 273 Kelvin
1 atm = 760 mmHg = 101.3 kPa
@ STP 1 mole gas = 22.4 L
Gas Laws:
Boyle’s Law: P1V1=P2V2
Charles’ Law: V1/T1= V2/T2
Lussac’s Law: P1/T1=P2/T2
Combined Gas law: P1V1/T1=P2V2/T2
Ideal Gas Law: PV=nRT