exam 2 Flashcards
chemical bonds
forces that hold atoms together
octet rule
atoms tend to bond to fill their valence orbital
- atoms seek to attain a “noble gas electron configuration”
2 types of covalent bonds
polar covalent: sharing of electrons is not as equal
non polar covalent: equal sharing of electrons
“which one has more ionic character?” means
which one has the greater electronegativity difference (more ionic)
ionic bonds
electron transfer where you have a lot of electronegativity
THEY SELFISH
ionic compounds are comprised of ___________
oppositely charged ions. they form when metals react with nonmetals.
*transfer of electrons from metal to nonmetal
6 properties of ionic compounds
- Contain a metal and nonmetal (large electronegativity difference)
- Bond strength (very strong bonds minus cellular conditions)
- **High melting and boiling points
** (b/c lots of negative & positive interactions – very stable & don’t want to be broken apart) - Crystalline structure
- Solubility in water (very soluble in water – why they’re weak in cellular conditions b/c cells are full of water)
- Conductivity (very low conductivity in solid state, however very soluble in H2O so high conductivity in aqueous solution)
bond length
distance b/w 2 atoms where the energy is lowest
covalent bonds form when ___________
electrons are shared between bonds
nonpolar covalent bonds vs. polar covalent bonds
nonpolar covalent: equal sharing between atoms in a molecule
polar covalent: unequal sharing between atoms in a molecule
5 properties of molecular (covalent) compounds
- Largely form when nonmetals react with other nonmetals
- Weak bond strength (weak bonds – broken much easier than ionic bonds)
- Often have low melting and boiling points
- Low solubility in water
- Low conductivity (b/c no ions are floating around)
How are binary ionic compounds named? (one metal and one nonmetal)
binary ionic compounds: contain a cation and anion
- name of the cation (metal) is named first and unchanged
- name of the anion (nonmetal) contains the suffix -ide
examples:
CsF (cesium fluoride)
MgO (magnesium oxide)
Ca3P2 (calcium phosphide)
Al4C3 (aluminum carbide)
Na2O (sodium oxide)
aming binary compounds (transition metals) + exceptions
- for atoms that can form different types of ions
- use ROMAN NUMERALS for charge
exceptions: Zn2+ , Ag+, Cd2+ (don’t need to use roman numerals for these elements)
how to name compounds using polyatomic ions
- special names that need to be memorized
use parentheses when you have more than 1 polyatomic
naming ionic hydrates
- naming is the same as ionic but has the extra aspect of hydrate in it
- naming hydrates: use Greek prefixes and the word “hydrate”
how to name binary covalent compounds
binary covalent compounds: contain two nonmetals
- first element is named
- second element contains suffix -“ide”
- prefixes used to denote number of atoms present for elements
mono never used for first element but used for the second element
ex. NO2 - nitrogen dioxide, IBr7- iodine heptabromide
what are acids
substances that release H+ when dissolved in water
how to name acids
- if anion ends in -ide, add prefix hydro- and suffix -ic
ex. HCL - hydrochloric acid - if anion ends in -ate, don’t use hydro- and add suffix -ic
ex. H2SO4 - sulfuric acid - if anion ends in -ite, -ite is replaced with -ous
ex. HNO2 - nitrous acid
this only applies for when they are in water, because in gaseous states, they are named differently
how are Lewis structures drawn for molecules
using only valence electrons, create dots and bonds on the name of the element
difference in Lewis structures for ionic compounds and molecules (covalent)
They are only helpful for covalent molecules, they don’t do much to show anything for ionic molecules
Octet and Duet Rule
octet rule: main group element want to obtain 8 valence electrons
g*duet rule**: hydrogen atoms may have no more than 2 electrons in their valence shells
exception of octet rule: molecules with odd numbers of electrons
atoms in molecules with an odd number of electrons cannot form stable octets
- atom might not enough to form 8 (fine)
- more electronegative atom will have more electrons
exception 2 to octet rule: electron-deficient molecules
molecules with not enough electrons will have central atoms with fewer than 8 electrons
exception 3 to octet rule: hypervalent molecules know this
- extra electrons can be placed on the central atom
- only for elements in period 3 and beyond
- often happens for sulfur and phosphorus
formal charge and formula
formal charge: charge assigned to an atom in a molecule (hypothetically: made up #’s)
FC = V - N - B )
V - valence electrons
N - non bonding electrons (one pairs)
B - # of bonds
2 predictions for formal charges
- atoms in molecules prefer charges close to zero
- any negative formal charges should reside on the more electronegative atoms
guiding principles- don’t apply everywhere but usually!
what is resonance and how are resonance structures drawn
resonance: occurs when multiple valid Lewis structures can be written for a single molecule
- all atoms must be connected in the same order, only double bond changes
- resonating structures have parenthesis around them and are connected by double headed arrows
resonance structures: superposition and major contributors
superposition: put them all on top and the average almost of the resonance structures
major contributor: contributes most to what the molecule looks like in nature, the one with the formal charges all closest to 0
electron arrangement vs. molecular structure
electron arrangement: describes general structure of the molecule (Lewis diagrams)
molecular structure: determined by the location of the atoms (VESPER)
AXnEm notation (for central atom)
A - central atom
X - number of bonds
E - number of lone pairs
ex: AX2E2 means central atom has 2 bonds and 2 sets of lone pairs
For VESPER model, multiple bonds on central atoms count as _________ electron group
one
For VESPER, molecules with multiple central atoms, ________
consider each center individually (one center could be tetrahedral and the trigonal pyramidal)
valence bond theory
says that bonding is basically the overlap of orbitals
- strength of covalent bond depends on the extent of overlap of the orbitals involved, more overlap = stronger the bond
sigma and pi bonds (orbital overlap)
sigma bonds: covalent bonds formed by head on overlap of two orbitals (oriented linearly)
pi bonds: covalent bonds formed by the side-by-side overlap of two orbitals
single, double, and triple bonds in terms of pi and sigma bonds
- single bond: 1 sigma bond
- double bond: 1 sigma, 1 pi bond
- triple bond: 1 sigma, 2 pi bond
hybridization definition and how it works
refers to the process of mixing atomic orbitals to form hybrid orbitals used for bonding
- only exist in covalently bonded atoms + only used for central atoms
- double bonds and triple bonds count as 1
- electron lone pairs count as 1
2: sp
3: sp2
4: sp3
5: sp3d
6: sp3d2
atoms vs molecules vs formula units
atoms: smallest particle of an element
molecules: two or more atoms joined together (strictly used for covalent compounds)
formula units: molecules but used for ionic compounds
diatomic elements
elements that naturally exist as molecules (H2, N2, O2, F2, Cl2, Br2, I2)
isomers
compounds with the same chemical formula but different molecular structures
formula mass
sum of the average atomic masses of all atoms present in the substance
- units: atomic mass units (amu)
molecular formulas, empirical formulas, and structural formulas
molecular formula: long one with all of the amounts of elements
empirical formula: simplified version of molecular
structural formula: actual drawings of the compounds
molar mass
mass (in grams) of 1 mole of substance
- numerically equivalent to a substance’s formula mass with units of g/mol
how do you find an element’s percent composition?
grams of that element over total grams of compound x 100
how do you determine the empirical formula given the percent composition?
assume out of 100 grams and make each percent value into grams
how do you find the molecular formula given the empirical and the gfm of the molecular?
divide molecular over empirical and the value you get is what each atom is multiplied by
solvent vs solute +aqueous
solvent: medium in which components are dissolved
solute: what’s being dissolved
aqueous solutions contain substances dissolved in liquid water
molarity + formula
molarity is a unit of concentration
M = moles of solute/ liters of solution
dilution + formula
dilution: processing of decreasing the concentration of a solution by adding solvent
C1V1 = C2V2
- the 1 is always for initial and the 2’s are for final ones, don’t confuse them!!*
this formula is ONLY for dilution, not anything else
used for stock solutions: concentrated solutions which are used to make less concentrated solutions
mass-volume percentage + formula
mass-volume percentage: (also called m/v%), mixed unit
mass component/ volume solution x 100
units: g/mL or kg/L
ppm & ppb + solutions
ppm: parts per million
mass of solute/mass of solution x 10^6
ppb parts per billon
mass of solute/mass of solution x 10^9
precipitation reactions
involve the formation of a precipitate (insoluble substance) when 2 solutions are mixed
- have to use the solubility rules table
when writing chemical formulas do you have to include the states?
YES
complete ionic equations vs. net ionic equations
complete includes all ionic equations and net cancels out the spectator ions
how to write complete ionic equations?
balance all the elements first, the coefficients of the compounds becomes coefficients of ions
- subscripts also turn into coefficients
acid base reactions and their products
Acid-base reactions: involve transfer of H+ from one species to another
acids dissolve in water to produce hydronium ions (H3O+)
bases dissolve in water to produce hydroxide ions (OH-)
neutralization reactions + products
occur when an acid and base react to form salt and water
strong & weak acids and bases + list of strong acids+bases
strong acids and bases: completely dissociate when dissolved in water
- doesn’t have to do with how dangerous it is to you
*7 strong acids**:
- HCl
- HBr
- HI
- HNO3
- HClO3
- HClO4
- H2SO4
strong bases:
- hydroxides of group I and II metal cations
oxidation-reduction (redox) reactions
involve transfer (or rearrangement of electrons)
electron acceptors = oxidizing agents
electron donors = reducing agents
the sum of oxidation numbers for all molecules or ions must equal ____________
the charge of the molecule or ion
oxidation states allow us to keep track of electrons during reactions.
what are the 5 basic rules of assigning oxidation states?
- an element by itself is 0
- ex. Na (s) and O2 (g)
- monoatomic ion is the same as its charge
- ex. Na+ or Cl-
- hydrogen is +1 in its covalent compound
- ex. H2O, HCl
- Oxygen is 2- for most compounds (*exception: peroxides (O2 2-) in which oxygen is -1 to maintain the overall charge of 2-)
- ex. H2O, CO2
- F is -1 always
- ex. HF
what is a limiting reactant?
reactant that runs out first, making further reaction impossible
how is the gfm of one mole of a diatomic calculated?
has to be multiplied by 2 bc technically 2 of that molecule
theoretical & percent yield + formula for percent yield
theoretical yield: amount of product formed when the limiting reactant is completely consumed
percent yield: percentage of the theoretical yield
equation = PAT
% yield = actual yield/ theoretical yield x 100
quantitative analysis + titration + gravimetric analysis definitions
quantitative analysis: determination of the amount or concentration of a substance in a sample
titration: adding volume of known concentration to the solution being analyzed, completion indicated by noticeable change in the solution like a color change
gravimetric analysis: involves a change (physical or chemical) during which an analyze is separated from other components of a sample
- usually forms a precipitate, more of an umbrella term
