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Intro to Chemistry > Finals > Flashcards

Finals Flashcards

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1
Q

chemistry

A

the science that deals with matter and the changes that matter undergoes

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2
Q

matter

A

the materials of the universe

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3
Q

chemical change

A

where one or more substances become different substances

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4
Q

scientific method

A
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5
Q

observation

A

qualitative

quantitative

can be witnessed/recorded

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6
Q

qualitative observation

A

doesn’t involve numbers

color, odor, appearance

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7
Q

quantitative observation

A

involves a # (mass, volume) & units

measurement

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8
Q

hypothesis

A

possible explanation for the observations

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9
Q

experiment

A

test hypothesis

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10
Q

theory/model

A

why?

once u have hypothesis that agrees w/ observations

a set of tested hypotheses that explains some part of nature

an interpretation of the behavior of nature

changes as more info becomes available

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11
Q

law of conservation of mass

A

the total mass of materials involved is the same before and after a chemical change

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12
Q

natural law

A

generally observed behavior as a statement

often see that same behavior applies to many diff systems

Ex: law of conservation of mass

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13
Q

law

A

what?

a summary of observed behavior

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14
Q

scientific notation

A

method used to make writing very large/very small #’s more compact and easier to use

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15
Q

units

A

define the scale of measurement being used

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16
Q

English system

A

mass - lb

length - ft

time - s

temp - F

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17
Q

SI System

A

mass - kg

length - m

time - s

temp - K

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18
Q

metric prefixes

A

Good Mornings kan definitely call me my name

G, giga, 109

M, mega, 106

k, kilo, 103

d, deci, 10-1

c, centi, 10-2

m, milli, 10-3

µ, micro, 10-6

n, nano, 10-9

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19
Q

volume

A

the amount of space that an obj occupies

metric - liter (L)

SI - m3

in lab, milliliter (mL) ↔ cm3 (cc)

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20
Q

mass

A

quantity of matter in an obj

SI - kg

in lab, gram(g)

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21
Q

certain digits

A

always the same

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22
Q

uncertain digits

A

estimated & may vary

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23
Q

significant figures

A

’s recorded in a measurement

determined by uncertainty of the measurement

do: nonzero, captive 0’s, trailing 0’s (if decimal)
infinite: exact & counted #’s

don’t: conversion factors, sci notation, leading 0’s

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24
Q

rounding off

A

carry all digits until final calculation

if first insignificant digit is 5+, up; -5 down

NOT sequential

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25
Q

sig fig multiplication/division

A

use # of sig figs in limiting term

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26
Q

limiting term

A

the measurement w/ the smallest # of sig figs or decimal places

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27
Q

sig fig addition/subtraction

A

smallest # of decimal places

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28
Q

accuracy

A

how close the measurement is to the true value

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29
Q

precision

A

a measure of how close the measurements are to each other

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30
Q

equivalence statement

A

defines the relationship b/w diff units

Ex: 1 kg = 2.205 lbs

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31
Q

conversion factor

A

a ratio of the 2 parts of an equivalence statement that relates the 2 units

Ex: 1kg/2.205lbs

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32
Q

dimensional analysis

A

the process of converting from 1 unit to another

choose conversion factor that cancels the units u don’t want

when cubing/squaring, make sure to cube/square the WHOLE conversion factor

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33
Q

Kelvin

A

absolute temp scale

doesn’t use degree notation, just K

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34
Q

extensive property

A

depends on quantity of sample measured

mass, volume

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35
Q

intensive property

A

independent of sample size

prop’s often characteristic of substance being measured

Ex: density, temp, melting & boiling pts

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36
Q

density

A

D = m/v

g/cm3 or g/mL

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37
Q

percent error

A

|accepted value - experimental value|

__________________________________

accepted value

x 100

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38
Q

properties of matter

A

physical

chemical

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39
Q

physical properties

A

properties that don’t involve substances changing into another substance

color, odor, taste, feel, density, melting & boiling pts, temp

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40
Q

chemical properties

A

prop’s that involve substances changing into another substance

chemical reaction

Ex: sugar ferments → alcohol, platinum doesn’t react w/ oxygen @ room temp, copper sheets on statue of liberty have green coating

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41
Q

matter

A

has volume & mass

three states: solid, liquid, gas

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42
Q

volume

A

the amount of space that an object occupies

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43
Q

mass

A

the amount of matter that an object contains

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44
Q

solid

A

rigid

fixed shape & volume

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45
Q

liquid

A

has definite volume but takes shape of container

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46
Q

gas

A

no fixed volume/shape

takes shape & volume of container

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47
Q

chemical reaction

A

1+ substances r changed into other substances

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48
Q

physical change

A

involves a change in 1+ phys prop’s, but no change in the fundamental components that make up the substance

most common = changes of state

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49
Q

changes of state

A

solid → liquid = melting

liquid → solid = freezing

liquid → gas = evaporation

gas → liquid = condensation

gas → solid = sublimation (dry ice)

solid → gas = deposition

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50
Q

chemical change

reaction

A

involves a change in the fundamentel components of the substance

a given substance changes into a diff substance/substances

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51
Q

element

A

a substance that can’t be broken down into other substances by chemical methods

microscopic form - sometimes used to mean a single atom of that element

macroscopic form - other times used to mean a sample large enough to weigh on a balance

generic form - when we say human body contains sodium, doesn’t mean elemental sodium is present, rather atoms of some form of sodium

118 elements, 88 of which occur naturally

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52
Q

compound

A

a substance composed of a given combo of elements that can be broken down into those elements by chem methods

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53
Q

mixture

A

a combo of substances in varying proportions

Ex: salt water

homogenous/heterogeneous

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54
Q

homogeneous mixture

solution

A

uniform composition

Ex: mixed salt water

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55
Q

heterogeneous mixture

A

non-uniform composition

Ex: choc chip cookie, sand & water

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56
Q

pure substance

A

will always have same composition

element/compound

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57
Q

organization of matter

A
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58
Q

separation of mixtures

A

physical changes

distillation

filtration

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59
Q

distillation

A

liquid → gas → liquid

separates liquids based on boiling pt by condensing vapor

used to recover liquids

Ex: salt water: solution heated to vaporize(boil) water. water vapor cooled so that condenses back to liquid state & all liquid is collected. after all water vaporized from original sample, pure sodium chloride remains.

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60
Q

filtration

A

separates solids from liquids

used to recover solids or liquids

Ex: mixture of salt(NaCl) & sand. sand = insoluble in water. add water & dissolve salt. filter so that salt solution passes thru & sand remains on filter. water then evaporated from salt.

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61
Q

evaporation

A

separates solids from liquids

used to recover solids

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62
Q

distilled water

A

water that has been evaporated & condensed to remove impurities

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63
Q

reagent

A

a substance or compound that is added to a system in order to bring about a chemical reaction, or added to see if a reaction occurs

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64
Q

precipitate

A

the formation of a solid in a solution during a chemical reaction

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65
Q

percent recovery

A

new total mass

______________

original total mass

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66
Q

mass percent of ___ in mixture

A

mass of recovered __

__________________

total mass of recovered solids

x 100

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67
Q

M

A

molarity

measure of concentration (moles/liter)

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68
Q

chemical symbols

A

used as abbreviations for element names

3 letters = unknown

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69
Q

Dalton’s Atomic Theory

A
  • most natural materials are mixtures of pure substances
  • pure substances are either elements or compounds
  • law of constant composition
  • elements are made of atoms
  • all atoms of a given element r identical
  • the atoms of a given element r diff from those of any other element
  • atoms of 1 element can combine with atoms of other elements to form compounds
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70
Q

compounds

A

a distinct substance that is composed of the atoms of 2+ elements and always contains exactly the same relative masses of those elements

can be broken down into elements by chem methods

always has the same relative #’s & types of atoms

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71
Q

law of constant composition

A

a given compound always contains the same proportion (by mass) of the elements

Ex: water always contains 8g of oxygen for each gram of hydrogen

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72
Q

atoms

A

tiny particles of which elements are made

indivisible in chem processes

not created/destroyed in chem reactions. reaction only change the way the atoms r grouped together

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73
Q

chemical formula

A

expresses the type of atoms & #’s of each atom in a given compound

table salt = NaCl

water = H2O

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74
Q

chemical name

A

table salt = sodium chloride

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75
Q

electron

A

a negatively charged particle

located outside of nucleus in a negatively charged “cloud”

most of volume of atom

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76
Q

proton

A

a positively charged particle, same size of charge as an electron, but positive

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77
Q

neutron

A

a neutral particle w/ a mass relatively close to that of a proton

78
Q

nucleus

A

small dense center of atom

protons & neutrons

most of the mass of the atom

79
Q

why do diff atoms have diff chem prop’s?

A

the # and arrangement of the electrons

the space in which the electrons move accounts for most of the atomic volume. the electrons r the parts of the atoms that intermingle when atoms combine to form molecules. the # of electrons a given atom possesses greatly affects the way it can interact w/ other atoms

80
Q

atomic number

A

the identity of an element is determined by this

the # of protons

81
Q

isotopes

A

(of an element)

have the same # of protons, & therefore the same atomic #, but diff #’s of neutrons

82
Q

mass number

A

the total # of protons & neutrons in an atom

used in naming to identify isotopes, such as Carbon-14 & Carbon-12

83
Q

practical purpose of isotopes

A

Iodine-123 = the isotope of choice for nuclear medicine imaging of the thyroid gland, which naturally accumulates all iodine isotopes

84
Q

metals

A

good conductors of heat & electricity

shiny

malleable

ductile

always form cations; tendency to lose electrons

85
Q

malleable

A

can be hammered into sheets

flattens/bends w/out shattering

86
Q

ductile

A

can be drawn into wires

87
Q

comparison of electron models

A
88
Q

nonmetals

A

good insulators

dull appearance

most r gases/easily vaporized solids & liquids

solids nonmetals r brittle

89
Q

insulator

A

absorbs electricity

90
Q

metalloids

A

exhibits some prop’s of both metals & nonmetals

91
Q

groups

A

all elements in a group have similar chem prop’s

92
Q

Group 1

A

Alkali Metals

93
Q

Group 2

A

Alkaline Earth Metals

94
Q

Groups 3-12

A

Transition Metals

95
Q

Group 17

A

Halogens

96
Q

Group 18

A

Noble Gases

97
Q

diatomic molecules

A

cannot stand alone

2 atoms

hydrogen, nitrogen, oxygen, flourine, chlorine, bromine, iodine

HINClBrOf

98
Q

ions

A

an equal # of protons & electrons gives an atom a net zero charge

adding or removing an electron from a neutral atom will create a charged ion

always formed by adding/removing electrons, not by changing the # of protons

99
Q

cation

A

positive ion

when atom loses 1+ electrons

Na → Na+ + e-

named using name of the parent atom

Ex: Na+ is called the sodium ion

100
Q

anion

A

negative ion

when an atom gains 1+ electrons

Cl + e- → Cl-

named by taking the root name of the atom & changing the ending (adding -ide)

101
Q

common anion name changes

A

chlorine - chloide ion

flourine - flouride

bromine - bromide

iodine - iodide

oxygen + 2 electrons = oxide

sulfur + 2 electrons = sulfide

102
Q

ion group charges

A

Group 1 metals → +1

Group 2 metals → +2

many Group 3-12 metals → multiple charges

Group 13 metals → +3

Group 16 atoms → -2

Group 17 atoms → -1

103
Q

ionic compounds

A

whenever a compound is formed b/w a metal & non-metals, it can be expected to contain ions

usually formed with metals + nonmetals; when they react, the metal atoms tend to lose 1+ electrons, which r gained by the atoms of the nonmetals; the reactions tend to form compounds that contain metal cations & nonmetal anions

chem compounds must have net charge of 0 → must be cations & anions present; the # of cations & anion must result in net charge of zero

usually strong electrolytes & can be expected to dissociate completely in solution

104
Q

rules for compound formula writing

A

the cation/metal/pos. is always written first

use subscripts to balance charges on compounds

105
Q

brittle

A

shatters/cracks into small pieces when struck

106
Q

strong electrolyte

A

a substance that separates into ions when dissolved in water

107
Q

dissociation

A

the process in which ionic ompounds separate into ions

108
Q

molecular compounds

A

usually non-electrolytes & don’t dissociate to form ions

resulting solutions don’t conduct electricity

109
Q

molecular acids

A

can partially/completely dissociate, depending on strength

110
Q

W

A

tungsten

74

111
Q

binary compounds

A

composed of 2 elements

two classes:

  1. metal & non-metal - metal forms only one cation
  2. metal & non-metal - the metal can form 2+ cations that have diff charges
  3. two non-metals
112
Q

Type I naming

A
  1. the cation is always named 1st and the anion 2nd
  2. a simple cation (obtained from a single atom) takes its name from the name of the element.
  3. a simple anion (obtained from a single atom) is named by taking the 1st part of the element name(the root) and adding -ide
113
Q

Type II naming

A
  1. cation always named 1st and the anion 2nd
  2. because the cation can assume more than one charge, the charge is specified by a Roman numeral in parentheses
114
Q

Type III naming

A
  1. 1st element in the formula is named first, and the full element name is used
  2. the 2nd element is named as though it were an anion
  3. prefixes are used to denote the numbers of atoms present
  4. the prefix mono- is never used for naming the 1st element
115
Q

polyatomic ions

A

charged entities composed of several atom bound together

name the cation first and then the anion. use Roman numerals if necessary.

116
Q

oxyanions

A

series of ions that contain a given element and different oxygen atoms

when there are 2 members in such a series, the one with the smaller amount of oxygen is called -ite and the one with the larger amount of oxygen = -ate. when there are more than two members of a series, hypo- (one less) and per- (one more) are used as prefixes

Ex: hypochlorite, chlorite, chlorate, perchlorate

117
Q

naming acids

A

when dissolved in water, certain molecules produce H+ ions → acids

an acid can have one/more H+ ions. the rules for naming acids depends on whether the anion contains oxygen

  • if the anion doesn’t contain oxygen, the acid is named with the prefix hydro- and the suffix -ic attached to the root name for the element. HCl is hydrochloric acid.
  • when the anion contains oxygen and ends in -ate, the suffix for the acid becomes -ic. H2SO4 is sulfuric acid
  • when the anion contains oxygen and ends in -its, the suffix for the acid becomes -ous. H2SO3 is sulfurous acid.
118
Q

common acids

A

HCl → hydrochloric acid

HC2H3O2 → acetic acid

H2SO4 → sulfuric acid

HNO3 → nitric acid

H3PO4 → phosphoric acid

119
Q

emission of energy by atoms

A

when atoms receive energy, they become excited. they can release the energy by emitting a light. the emitted energy is carried away by a photon

the energy of the photon corresponds exactly to the energy change of the emitting atom

high energy photons correspond to short wavelength light. low energy photons correspond to long wavelength light

the photons of red light have less energy than the photons of blue light cuz red light has a longer wavelength than blue light

120
Q

energy levels of hydrogen

A

when we study the photons of visible light emitted, we only see certain colors

only certain types of photons r produced

because only certain protons r emitted, only certain energy changes r occurring

so, hydrogen atoms must have certain discrete energy levels

we say the energy levels of H r quantized, that is, only certain values r allowed

energy levels of all atoms r quantized

121
Q

hydrogen orbitals

A

the probability map is called an orbital. the orbital shown in the pic is called the 1s orbital & describes the ground (lowest) state of energy for H

size of orbital is defined by a sphere that contains 90% of the total electron probability

122
Q

principal energy levels

A

discrete energy levels

designated by whole #’s symbolized by n; n can equal 1, 2, 3, 4,… level 1 corresponds to n = 1,etc.

energy of the level increases as value of n increases

describe size & shape. the s orbital is spherical. level 1 is smaller than level 2, which is smaller than level 3.

each principal energy level contains 1+ types of orbitals, called sublevels

123
Q

sublevels

A

the # of sublevels present in a given principal energy level equals n.

e.g. level 1 contains one sublevel (1s); level 2 contains 2 sublevels (2 types of orbitals), the 2s orbital and the three 2p orbitals; and so on

these r sumarized in the pic. the # of each type of orbital is shown in parentheses

n value always used to label orbitals of a given principal energy level & followed by letter that indicates the type/shape of the orbital

124
Q

orbital

A

can be empty or can contain 1 or 2 electrons, but never more than 2

if 2 electrons occupy the same orbital, must have opp spins

shape of an orbital doesn’t indicate the details of electron movement - merely indicates the prob distribution for an electron residing in that orbital

125
Q

spin

A

each electron appears to spin like a top on its axis

can only spin in 1 direction. we represent spin with up & down arrows

126
Q

Pauli Exclusion Principle

A

an atomic orbital can hold a max of 2 electrons & those 2 electrons must have opp spins

127
Q

electron configuration

A

principal energy level followed by sublevel; # of electrons in the orbital placed as superscript

Ex: 1s1

128
Q

orbital diagram

box diagram

A
129
Q

valence electrons

A

the electrons in the outermost(highest) principal energy level of an atom

these r the electrons involved in bonding of atoms to each other

the atoms of elements in the same group have the same # of electrons in a given type of orbital, except that the orbitals are in diff princ energy levels (except He, which is 1s2)

elements w/ same valence electron arrangement show very similar chem behavior

130
Q

orbital filling

A

in a principal energy level that has d orbitals, the s orbital from the next level fills before the d orbitals in the current level. that is, the (n + 1)s orbitals always fill before the nd orbitals

Ex: the 5s orbitals fill for rubidium & strontium before the 4d orbitals fill

131
Q

lanthanide series

A

after lanthanum, which has configuration [Xe]6s25d1

a group of 14 elements

corresponds to the filling of the seven 4f orbitals

132
Q

actinide series

A

after actinium, [Rn]7s26d1

14 elements

corresponds to the filling of seven 5f orbitals

133
Q

bond

A

a force that holds 2+ atoms together & makes them function as a unit

in water, the fundamental unit is the H-O-H molecule, which is held together by the two O-H bonds

134
Q

ionic compounds

A

formed when an atom that loses an electron relatively easily reacts w/ an atom that accepts an electron

when metal reacts w/ non-metal

resulting bonds = ionic bonds

electrons transferred

135
Q

covalent bond

A

when 2 similar atoms form a bond, the electrons r equaly attracted to the nuclei of the 2 atoms

electrons shared by nuclei

Ex: diatomic hydrogen H-H

136
Q

polar covalent bonds

A

b/w the extremes

atoms r not so diff that electrons r transferred, but diff enough that unequal sharing of the electrons results

137
Q

electronegativity

A

the unequal sharing of electrons b/w 2 atoms is described by this property

the relative ability of an atom in a molecule to attract shared electrons to itself

the higher the electronegativity value, the closer the shared electrons tend to be to that atom when it forms a bond

F has highest electronegativy & so always forms polar bonds

increasing electronegativity as goes right and up on periodic table

138
Q

polarity

A

depends on the diff b/w the electronegativity values of the atoms forming the bonds

if similar electronegativities, the electrons r shared almost equally & bond shows little polarity

if very diff electronegativities, very polar bond is formed

139
Q

stable electron configurations

A

representative (main-group) metals form ions by losing enough electrons to attain the configuration of the previous noble gas that occurs before the metal

nonmetals form ions by gaining enough electrons to attain the configuration of the next noble gas

when a non-metal and a Group 1, 2, or 3 metal react to form a binary ionic bond, the ions form so that the non-metal completes the valence-electron config of the next noble gas & the metal empties the valence orbitals to achieve the config. of the prev. noble gas

when 2 non-metals react to form a covalent bond, share electrons in a way that completes the valence-electron configuration of both atoms

140
Q

Lewis structures

A

bonding involves jsut the valence electrons. this structure is a representation of a molecule that shows how the valence electrons r arranged among the atoms in a molecule

H - duet rule

He - doesn’t form bonds cuz valence orbital already filled

2nd row nonmetals C thru F - octet rule

Ne - doesn’t form bonds cuz already has an octet of valence electrons

  1. obtain sum of the valence electrons from all of the atoms
  2. use 1 pair of electrons to form a bond b/w each pair of bound atoms. use line.
  3. arrange the remaining electrons to satisfy the duet rule for hydrogen & the octet rule for each 2nd-row element (may need to guess & check w/ double bonds, etc.)
141
Q

duet rule

A

H forms stable molecules where it shares only 2 electrons

142
Q

octet rule

A

2nd-row nonmetals C through F form stable molecules when r surrounded by enough electrons to fill the valence orbitals - that is, the one 2s and the three 2p orbitals

8 electrons required to fill these orbitals

143
Q

configuration tips

A

noble gas valences from top to bottom:

1s2

2s22p6

3s23p6

4s24p6

5s25p6

6s26p6

144
Q

evidence of a chemical reaction

A

color change

formation of a precipitate (solid)

formation of a gas (bubbles)

heat is produced (exothermic) or heat is absorbed (endothermic)

145
Q

chemical equation

A

reactants → products

146
Q

conservation of mass

A

in a chemical reaction, atoms r neither created nor destroyed

thr must be the same # of atoms on the reactant side of the equation as there are on the product side of the equation

147
Q

physical states

A

g - gas

l - liquid

s - solid

aq - aqueous

148
Q

what causes reactions?

A

precipitation reactions (driving force = formation of precipitate)

gas-forming reactions (driving force = formation of a gas)

acid-base reactions (driving force = formation of water)

transfer of electrons

149
Q

precipitation

A

formation of a solid

solid formed = precipitate

reaction = precipitation reaction

150
Q

predicting precipitates

A

insoluble

solid

151
Q

strong electrolyte

A

a substance that completely breaks apart into ions when dissolved in water

resulting solution readily conducts an electric current

Ba(NO3)2 and K2CrO4

152
Q

soluble solid

A

readily dissolves in water

153
Q

insoluble solid

slightly soluble solid

A

only a small amount of the solid dissolves in water

154
Q

ionic compound

A

all salts

when ionic compounds dissolve, the resulting solution contains ions

155
Q

predicting equations

A
  1. exchange anions & cations
  2. balance charges
  3. balance equation
  4. use solubility rule to find precipitates
156
Q

combination

synthesis reaction

A

two reactants combine to form a single product. the reactants may be elements or compounds

Zn(s) + I2(s) → ZnI2(s)

157
Q

decomposition

A

one reactant, a compound, breaks down to give 2+ products

2H2O2(aq) → 2H2O(l) + O2(g)

158
Q

single replacement

A

an element reacts with a compound and replaces one of the elements in the compound

metals replace hydrogen or other metals; nonmetals replace nonmetals

Zn(s) + 2HCl(aq) → H2(g) + ZnCl2(aq)

159
Q

double replacement

A

2 ionic compounds exchange ions to form new compounds

NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)

160
Q

combustion

A

a compound burns in the presence of oxygen, producing energy in the form of heat and light

the combustion of organic compounds produces carbon dioxide and water

C4H8(l) + 6O2(g) → 4CO2(g) + 4H2O(g)

161
Q

molecular equation

A

shows the overall reaction but not necessarily the actual forms of the reactants & products in solutions

balanced in charge & molecules

K2CrO4(aq) + Ba(NO3)2(aq) → BaCrO4(s) + 2KNO3(aq)

162
Q

complete ionic equation

A

represents all reactants & products that are strong electrolytes as ions

all reactants & products r included

2K+(aq) + CrO42-(aq) + Ba2+(aq) + 2NO3-(aq) → BaCrO4(s) + 2K+(aq) + 2NO3-(aq)

when writing ions, put charge on top right, and change subscript to coefficient (unless polyatomic - in which case, MAKE SURE the subscript is EXTRA)

163
Q

spectator ions

A

ions that don’t participate directly in a reaction

164
Q

net ionic equation

A

includes only those components that undergo a change

spectator ions r not included

Ba2+(aq) + CrO42-(aq) → BaCrO4(s) [Ba2+ and CrO42- both changed from aq to s]

165
Q

balancing tip

A

if 2 elements being combined have the same subscript, u can take it away and change it to a coefficient

166
Q

atomic mass

molecular weight (MV)

A

most elements occur in nature as a mixture of isotopes

average mass of an atom in an element, expressed in atomic mass units (amu) or grams/mole

this is one reason why atomic masses r not whole #’s - they are based on averages

the average atomic weight of an element can be calculated if the abundance of each isotope for that element is known

Ex: Chlorine = mixture of 2 isotopes

35Cl - 34.96885268 amu - 75.77% abundance

37Cl - 36.96590259 amu - 24.23% abundance

35Cl → (75.77/100) • 34.97 amu = 26.50 amu

37Cl → (24.23/100) • 36.97 amu = 8.95 amu

26.5amu + 8.95amu = 35.45 amu = average atomic mass for chlorine

167
Q

amu

A

atomic mass unit

= 1/12th of the mass of a 12C atom = 1/661 x 10-24 gram

168
Q

mole

A

number of atoms in 12.000g of 12C can be calculated

one atom 12C = 12.000 amu (by definition)

= 12.000amu x (1.661 x 10-24g/amu)

1 atom = 1.993 x 10-23g

number of atoms = 12.000g • (1 atom/1.993 x 10-23g) = 6.02 x 1023 atoms

  • the # of atoms of any element needed to equal its atomic mass in grams will always be 6.022 x 1023 atoms, a quantity known as the **mole **(also known as Avogadro’s number)
  • one mole equals the atomic mass in grams of an element
    • mass of 1 mole of 12C = 12.000g
    • mass of 1 mole of C = 12.011g
    • mass of 1 mole of Na = 22.990g
    • mass of 1 mole of H = 1.008g
    • mass of 1 mole of O = 15.999g
169
Q

formula mass

formula weight FW

A

total mass for all atoms in a compound

170
Q

molar mass

A

the mass (in grams) of 1 mol of the compound and is the sum of the masses of the component atoms

171
Q

calculating mass using amu

A

Calculate the mass, in amu, of a sample of aluminum that contains 75 atoms.

  1. 1 Al atom = 26.98 amu
  2. 75 Al atoms x (26.98amu/Al atom) = 2024amu
172
Q

calculating the number of atoms from the mass

A

Calculate the # of sodium atoms present in a sample that has a mass of 1172.49 amu

  1. 1 Na atom = 22.99 amu
  2. 1172.49amu x (1 Na/22.99amu) = 51.00 Na atoms
173
Q

calculating moles and number of atoms

A

Compute both the # of moles of atoms and the # of atoms in a 10.0-g sample of aluminum.

  1. 1 mol Al = 26.98 g Al
  2. 10.0g Al x (1mol Al/26.98g Al) = 0.371 mol Al
  3. 6.022 x 1023 Al atoms = 1 mol Al atoms
  4. 0.371 mol Al x (6.022 x 1023 Al atoms/1 mol Al) = 2.23 x 1023 Al atoms
174
Q

calculating the number of atoms

A

How many silicon(Si) atoms are present in 5.68mg? The average atomic mass for Si is 28.09 amu

  1. 1g = 1000mg
  2. 5.68mg Si x (1g Si/1000mg Si) = 5.68 x 10-3 g Si
  3. 1 mol Si atoms = 28.09 g Si
  4. 5.68 x 10-3 g Si x (1 mol Si/28.09 g Si) = 2.02 x 10-4 mol Si
  5. 1 mol = 6.022 x 1023
  6. 2.02 x 10-4 mol Si x (6.022 x 1023 atoms/1 mol Si) = 1.22 x 1020 Si atoms
175
Q

acid

A

proton (H+) donor

176
Q

base

A

a proton acceptor

177
Q

acids and bases

A

the general reaction that occurs when an acid is dissolved in water can best be rep-ed as an aicid donating a proton to a water molecule to form a new acid (the conjugate acid) and a new base (the conjugate base)

HA(aq)(acid) + H2O(l)(base) → H3O+(aq)(conj acid) + A-(aq)(conj base)

this model emphasizes the role of the water molecule in pulling the proton from the acid

178
Q

conjugate acid

A

formed when the proton is transferred to the base

179
Q

conjugate acid-base pair

A

2 substances related to each other by the donating and accepting of a single proton

HA(aq)(acid) + H2O(l)(base) → H3O+(aq)(conj acid) + A-(aq)(conj base)

↑↑↑there are 2 acid-base pairs: HA(acid) and A-(base) and H2O(base) and H3O+(acid)

180
Q

more acids and bases

A

HCl(aq)(acid) + H2O(l)(base) → H3O+(aq)(conj acid) + Cl-(aq)(conj base)

  • in this case, HCl is the acid that loses an H+ ion to form Cl-, its conjugate base
  • H2O (acting as a base) gains an H+ ion to form H3O+ (the conj acid)
  • H3O+ is called the hydronium ion
181
Q

which are acid-base pairs?

A
  1. HF, F-→ yes. lost H+
  2. NH4+, NH3 → yes. lost H+
  3. HCl, H2O → no

1 & 2 are conjugate acid-base pairs cuz the 2 species differ by one H+

182
Q

reverse reaction

A

the conj acid & conj base can also react w/ each other to reform the parent molecule and water. a reaction can occur in both directions (other = forward reaction)

the products in the forward reaction = the reactants in the reverse reaction

we use double arrows to rep reactions that occur in both directions

183
Q

strong acid

A

HA(aq) + H20(l) ↔ H3O+(aq) + A-(aq)

this rep’s a competition for the H+ ion b/w the H2O in the forward reaction and the A- in the reverse direction (either H+ or A- will go with H2O)

if the H2O has a higher attraction for H+ compared to A-, then the solution will contain mostly H3O+ and A-. The forward reaction dominates and the acid is completely ionized or dissociated.

has relatively weak conjugate base

184
Q

weak acid

A

if the H2O has a higher attraction for A-compared to H+ then the solution will contain mostly HA and H2O. the reverse reaction dominates and most of the acid remains as HA.

relatively strong conj base

Ex: acetic acid. most of the acid remains intact when dissolved in water.

185
Q

diprotic acid

A

contributes 2 H+ ions when dissolved in water

Ex: sulfuric acid, H2SO4

186
Q

amphoteric substance

A

can act as an acid or a base

H2O

in the ionization of water, a proton transfers from one molecule of water to the other, producing a hydroxide ion and a hyronium ion

H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq)

in this reaction, 1 water molecule acts as an acid by donating a proton. the other acts as a base by accepting a proton

the forward reaction doesn’t occur to a great extent, meaning that in pure water, very little amounts of hydronium and hydroxide ions exist.

187
Q

concentration

A

[] used to denote this

[H3O+] = [OH-] = 1.0 x 10-7 M in water

[H3O+] can also be written as [H+]

(volume of acid)(concentration of acid) = (volume of base)(concentration of base)

Va • Ma = Vb • Mb

measured in units of solute/liter of solution -or- molarity

188
Q

ion-product constant

A

the product of [H+] [OH-] = (1.0 x 10-7(1.0 x 10-7) = 1.0 x 10-4 and is a constant

noted as Kw

Kw = [H+] [OH-] = 1.0 x 10-4

** because the product is constant, if [H+] increases, [OH-] must decrease. If [H+] decreases, [OH-] must increase

189
Q

possible acid situations

A

in a solution, there are 3 possible situations

  1. adding an acid to water, increasing [H+]

[H+] > [OH-] → acidic solution

  1. adding a base to water, increasing [OH-]

[H+] < [OH-] → basic solution

  1. having equal amounts of acid and base

[H+] = [OH-] → neutral solution

190
Q

pH

A

a convenient way to express the acidity of a solution

pH = -log[H+]

1-6 = acidic

7 = neutral

8-14 = basic

Intro to Chemistry (6 decks)

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