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CHEM 101 > exam 2 > Flashcards

exam 2 Flashcards

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1
Q

chemical bonds

A

forces that hold atoms together

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2
Q

octet rule

A

atoms tend to bond to fill their valence orbital
- atoms seek to attain a “noble gas electron configuration”

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3
Q

2 types of covalent bonds

A

polar covalent: sharing of electrons is not as equal

non polar covalent: equal sharing of electrons

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4
Q

“which one has more ionic character?” means

A

which one has the greater electronegativity difference (more ionic)

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5
Q

ionic bonds

A

electron transfer where you have a lot of electronegativity

THEY SELFISH

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6
Q

ionic compounds are comprised of ___________

A

oppositely charged ions. they form when metals react with nonmetals.

*transfer of electrons from metal to nonmetal

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7
Q

6 properties of ionic compounds

A
  1. Contain a metal and nonmetal (large electronegativity difference)
  2. Bond strength (very strong bonds minus cellular conditions)
  3. **High melting and boiling points
    ** (b/c lots of negative & positive interactions – very stable & don’t want to be broken apart)
  4. Crystalline structure
  5. Solubility in water (very soluble in water – why they’re weak in cellular conditions b/c cells are full of water)
  6. Conductivity (very low conductivity in solid state, however very soluble in H2O so high conductivity in aqueous solution)
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8
Q

bond length

A

distance b/w 2 atoms where the energy is lowest

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9
Q

covalent bonds form when ___________

A

electrons are shared between bonds

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10
Q

nonpolar covalent bonds vs. polar covalent bonds

A

nonpolar covalent: equal sharing between atoms in a molecule

polar covalent: unequal sharing between atoms in a molecule

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11
Q

5 properties of molecular (covalent) compounds

A
  1. Largely form when nonmetals react with other nonmetals
  2. Weak bond strength (weak bonds – broken much easier than ionic bonds)
  3. Often have low melting and boiling points
  4. Low solubility in water
  5. Low conductivity (b/c no ions are floating around)
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12
Q

How are binary ionic compounds named? (one metal and one nonmetal)

A

binary ionic compounds: contain a cation and anion

  1. name of the cation (metal) is named first and unchanged
  2. 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)

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13
Q

aming binary compounds (transition metals) + exceptions

A
  • 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)

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14
Q

how to name compounds using polyatomic ions

A
  • special names that need to be memorized

use parentheses when you have more than 1 polyatomic

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15
Q

naming ionic hydrates

A
  • naming is the same as ionic but has the extra aspect of hydrate in it
  • naming hydrates: use Greek prefixes and the word “hydrate”
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16
Q

how to name binary covalent compounds

A

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

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17
Q

what are acids

A

substances that release H+ when dissolved in water

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18
Q

how to name acids

A
  • 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

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19
Q

how are Lewis structures drawn for molecules

A

using only valence electrons, create dots and bonds on the name of the element

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20
Q

difference in Lewis structures for ionic compounds and molecules (covalent)

A

They are only helpful for covalent molecules, they don’t do much to show anything for ionic molecules

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21
Q

Octet and Duet Rule

A

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

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22
Q

exception of octet rule: molecules with odd numbers of electrons

A

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
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23
Q

exception 2 to octet rule: electron-deficient molecules

A

molecules with not enough electrons will have central atoms with fewer than 8 electrons

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24
Q

exception 3 to octet rule: hypervalent molecules know this

A
  • extra electrons can be placed on the central atom
  • only for elements in period 3 and beyond
  • often happens for sulfur and phosphorus
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25
Q

formal charge and formula

A

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

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26
Q

2 predictions for formal charges

A
  • 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!

27
Q

what is resonance and how are resonance structures drawn

A

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
28
Q

resonance structures: superposition and major contributors

A

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

29
Q

electron arrangement vs. molecular structure

A

electron arrangement: describes general structure of the molecule (Lewis diagrams)

molecular structure: determined by the location of the atoms (VESPER)

30
Q

AXnEm notation (for central atom)

A

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

31
Q

For VESPER model, multiple bonds on central atoms count as _________ electron group

A

one

32
Q

For VESPER, molecules with multiple central atoms, ________

A

consider each center individually (one center could be tetrahedral and the trigonal pyramidal)

33
Q

valence bond theory

A

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
34
Q

sigma and pi bonds (orbital overlap)

A

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

35
Q

single, double, and triple bonds in terms of pi and sigma bonds

A
  • single bond: 1 sigma bond
  • double bond: 1 sigma, 1 pi bond
  • triple bond: 1 sigma, 2 pi bond
36
Q

hybridization definition and how it works

A

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

37
Q

atoms vs molecules vs formula units

A

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

38
Q

diatomic elements

A

elements that naturally exist as molecules (H2, N2, O2, F2, Cl2, Br2, I2)

39
Q

isomers

A

compounds with the same chemical formula but different molecular structures

40
Q

formula mass

A

sum of the average atomic masses of all atoms present in the substance

  • units: atomic mass units (amu)
41
Q

molecular formulas, empirical formulas, and structural formulas

A

molecular formula: long one with all of the amounts of elements

empirical formula: simplified version of molecular

structural formula: actual drawings of the compounds

42
Q

molar mass

A

mass (in grams) of 1 mole of substance

  • numerically equivalent to a substance’s formula mass with units of g/mol
43
Q

how do you find an element’s percent composition?

A

grams of that element over total grams of compound x 100

44
Q

how do you determine the empirical formula given the percent composition?

A

assume out of 100 grams and make each percent value into grams

45
Q

how do you find the molecular formula given the empirical and the gfm of the molecular?

A

divide molecular over empirical and the value you get is what each atom is multiplied by

46
Q

solvent vs solute +aqueous

A

solvent: medium in which components are dissolved

solute: what’s being dissolved

aqueous solutions contain substances dissolved in liquid water

47
Q

molarity + formula

A

molarity is a unit of concentration

M = moles of solute/ liters of solution

48
Q

dilution + formula

A

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

49
Q

mass-volume percentage + formula

A

mass-volume percentage: (also called m/v%), mixed unit

mass component/ volume solution x 100

units: g/mL or kg/L

50
Q

ppm & ppb + solutions

A

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

51
Q

precipitation reactions

A

involve the formation of a precipitate (insoluble substance) when 2 solutions are mixed

  • have to use the solubility rules table
52
Q

when writing chemical formulas do you have to include the states?

A

YES

53
Q

complete ionic equations vs. net ionic equations

A

complete includes all ionic equations and net cancels out the spectator ions

54
Q

how to write complete ionic equations?

A

balance all the elements first, the coefficients of the compounds becomes coefficients of ions

  • subscripts also turn into coefficients
55
Q

acid base reactions and their products

A

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-)

56
Q

neutralization reactions + products

A

occur when an acid and base react to form salt and water

57
Q

strong & weak acids and bases + list of strong acids+bases

A

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

58
Q

oxidation-reduction (redox) reactions

A

involve transfer (or rearrangement of electrons)

electron acceptors = oxidizing agents

electron donors = reducing agents

59
Q

the sum of oxidation numbers for all molecules or ions must equal ____________

A

the charge of the molecule or ion

60
Q

oxidation states allow us to keep track of electrons during reactions.

what are the 5 basic rules of assigning oxidation states?

A
  1. an element by itself is 0
    • ex. Na (s) and O2 (g)
  2. monoatomic ion is the same as its charge
    • ex. Na+ or Cl-
  3. hydrogen is +1 in its covalent compound
    • ex. H2O, HCl
  4. 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
  5. F is -1 always
    • ex. HF
61
Q

what is a limiting reactant?

A

reactant that runs out first, making further reaction impossible

62
Q

how is the gfm of one mole of a diatomic calculated?

A

has to be multiplied by 2 bc technically 2 of that molecule

63
Q

theoretical & percent yield + formula for percent yield

A

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

64
Q

quantitative analysis + titration + gravimetric analysis definitions

A

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

CHEM 101 (6 decks)

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