Finals Flashcards
chemistry
the science that deals with matter and the changes that matter undergoes
matter
the materials of the universe
chemical change
where one or more substances become different substances
scientific method
observation
qualitative
quantitative
can be witnessed/recorded
qualitative observation
doesn’t involve numbers
color, odor, appearance
quantitative observation
involves a # (mass, volume) & units
measurement
hypothesis
possible explanation for the observations
experiment
test hypothesis
theory/model
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
law of conservation of mass
the total mass of materials involved is the same before and after a chemical change
natural law
generally observed behavior as a statement
often see that same behavior applies to many diff systems
Ex: law of conservation of mass
law
what?
a summary of observed behavior
scientific notation
method used to make writing very large/very small #’s more compact and easier to use
units
define the scale of measurement being used
English system
mass - lb
length - ft
time - s
temp - F
SI System
mass - kg
length - m
time - s
temp - K
metric prefixes
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
volume
the amount of space that an obj occupies
metric - liter (L)
SI - m3
in lab, milliliter (mL) ↔ cm3 (cc)
mass
quantity of matter in an obj
SI - kg
in lab, gram(g)
certain digits
always the same
uncertain digits
estimated & may vary
significant figures
’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
rounding off
carry all digits until final calculation
if first insignificant digit is 5+, up; -5 down
NOT sequential
sig fig multiplication/division
use # of sig figs in limiting term
limiting term
the measurement w/ the smallest # of sig figs or decimal places
sig fig addition/subtraction
smallest # of decimal places
accuracy
how close the measurement is to the true value
precision
a measure of how close the measurements are to each other
equivalence statement
defines the relationship b/w diff units
Ex: 1 kg = 2.205 lbs
conversion factor
a ratio of the 2 parts of an equivalence statement that relates the 2 units
Ex: 1kg/2.205lbs
dimensional analysis
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
Kelvin
absolute temp scale
doesn’t use degree notation, just K
extensive property
depends on quantity of sample measured
mass, volume
intensive property
independent of sample size
prop’s often characteristic of substance being measured
Ex: density, temp, melting & boiling pts
density
D = m/v
g/cm3 or g/mL
percent error
|accepted value - experimental value|
__________________________________
accepted value
x 100
properties of matter
physical
chemical
physical properties
properties that don’t involve substances changing into another substance
color, odor, taste, feel, density, melting & boiling pts, temp
chemical properties
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
matter
has volume & mass
three states: solid, liquid, gas
volume
the amount of space that an object occupies
mass
the amount of matter that an object contains
solid
rigid
fixed shape & volume
liquid
has definite volume but takes shape of container
gas
no fixed volume/shape
takes shape & volume of container
chemical reaction
1+ substances r changed into other substances
physical change
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
changes of state
solid → liquid = melting
liquid → solid = freezing
liquid → gas = evaporation
gas → liquid = condensation
gas → solid = sublimation (dry ice)
solid → gas = deposition
chemical change
reaction
involves a change in the fundamentel components of the substance
a given substance changes into a diff substance/substances
element
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
compound
a substance composed of a given combo of elements that can be broken down into those elements by chem methods
mixture
a combo of substances in varying proportions
Ex: salt water
homogenous/heterogeneous
homogeneous mixture
solution
uniform composition
Ex: mixed salt water
heterogeneous mixture
non-uniform composition
Ex: choc chip cookie, sand & water
pure substance
will always have same composition
element/compound
organization of matter
separation of mixtures
physical changes
distillation
filtration
distillation
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.
filtration
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.
evaporation
separates solids from liquids
used to recover solids
distilled water
water that has been evaporated & condensed to remove impurities
reagent
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
precipitate
the formation of a solid in a solution during a chemical reaction
percent recovery
new total mass
______________
original total mass
mass percent of ___ in mixture
mass of recovered __
__________________
total mass of recovered solids
x 100
M
molarity
measure of concentration (moles/liter)
chemical symbols
used as abbreviations for element names
3 letters = unknown
Dalton’s Atomic Theory
- 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
compounds
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
law of constant composition
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
atoms
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
chemical formula
expresses the type of atoms & #’s of each atom in a given compound
table salt = NaCl
water = H2O
chemical name
table salt = sodium chloride
electron
a negatively charged particle
located outside of nucleus in a negatively charged “cloud”
most of volume of atom
proton
a positively charged particle, same size of charge as an electron, but positive
neutron
a neutral particle w/ a mass relatively close to that of a proton
nucleus
small dense center of atom
protons & neutrons
most of the mass of the atom
why do diff atoms have diff chem prop’s?
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
atomic number
the identity of an element is determined by this
the # of protons
isotopes
(of an element)
have the same # of protons, & therefore the same atomic #, but diff #’s of neutrons
mass number
the total # of protons & neutrons in an atom
used in naming to identify isotopes, such as Carbon-14 & Carbon-12
practical purpose of isotopes
Iodine-123 = the isotope of choice for nuclear medicine imaging of the thyroid gland, which naturally accumulates all iodine isotopes
metals
good conductors of heat & electricity
shiny
malleable
ductile
always form cations; tendency to lose electrons
malleable
can be hammered into sheets
flattens/bends w/out shattering
ductile
can be drawn into wires
comparison of electron models
nonmetals
good insulators
dull appearance
most r gases/easily vaporized solids & liquids
solids nonmetals r brittle
insulator
absorbs electricity
metalloids
exhibits some prop’s of both metals & nonmetals
groups
all elements in a group have similar chem prop’s
Group 1
Alkali Metals
Group 2
Alkaline Earth Metals
Groups 3-12
Transition Metals
Group 17
Halogens
Group 18
Noble Gases
diatomic molecules
cannot stand alone
2 atoms
hydrogen, nitrogen, oxygen, flourine, chlorine, bromine, iodine
HINClBrOf
ions
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
cation
positive ion
when atom loses 1+ electrons
Na → Na+ + e-
named using name of the parent atom
Ex: Na+ is called the sodium ion
anion
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)
common anion name changes
chlorine - chloide ion
flourine - flouride
bromine - bromide
iodine - iodide
oxygen + 2 electrons = oxide
sulfur + 2 electrons = sulfide
ion group charges
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
ionic compounds
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
rules for compound formula writing
the cation/metal/pos. is always written first
use subscripts to balance charges on compounds
brittle
shatters/cracks into small pieces when struck
strong electrolyte
a substance that separates into ions when dissolved in water
dissociation
the process in which ionic ompounds separate into ions
molecular compounds
usually non-electrolytes & don’t dissociate to form ions
resulting solutions don’t conduct electricity
molecular acids
can partially/completely dissociate, depending on strength
W
tungsten
74
binary compounds
composed of 2 elements
two classes:
- metal & non-metal - metal forms only one cation
- metal & non-metal - the metal can form 2+ cations that have diff charges
- two non-metals
Type I naming
- the cation is always named 1st and the anion 2nd
- a simple cation (obtained from a single atom) takes its name from the name of the element.
- a simple anion (obtained from a single atom) is named by taking the 1st part of the element name(the root) and adding -ide
Type II naming
- cation always named 1st and the anion 2nd
- because the cation can assume more than one charge, the charge is specified by a Roman numeral in parentheses
Type III naming
- 1st element in the formula is named first, and the full element name is used
- the 2nd element is named as though it were an anion
- prefixes are used to denote the numbers of atoms present
- the prefix mono- is never used for naming the 1st element
polyatomic ions
charged entities composed of several atom bound together
name the cation first and then the anion. use Roman numerals if necessary.
oxyanions
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
naming acids
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.
common acids
HCl → hydrochloric acid
HC2H3O2 → acetic acid
H2SO4 → sulfuric acid
HNO3 → nitric acid
H3PO4 → phosphoric acid
emission of energy by atoms
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
energy levels of hydrogen
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
hydrogen orbitals
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
principal energy levels
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
sublevels
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
orbital
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
spin
each electron appears to spin like a top on its axis
can only spin in 1 direction. we represent spin with up & down arrows
Pauli Exclusion Principle
an atomic orbital can hold a max of 2 electrons & those 2 electrons must have opp spins
electron configuration
principal energy level followed by sublevel; # of electrons in the orbital placed as superscript
Ex: 1s1
orbital diagram
box diagram
valence electrons
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
orbital filling
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
lanthanide series
after lanthanum, which has configuration [Xe]6s25d1
a group of 14 elements
corresponds to the filling of the seven 4f orbitals
actinide series
after actinium, [Rn]7s26d1
14 elements
corresponds to the filling of seven 5f orbitals
bond
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
ionic compounds
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
covalent bond
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
polar covalent bonds
b/w the extremes
atoms r not so diff that electrons r transferred, but diff enough that unequal sharing of the electrons results
electronegativity
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
polarity
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
stable electron configurations
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
Lewis structures
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
- obtain sum of the valence electrons from all of the atoms
- use 1 pair of electrons to form a bond b/w each pair of bound atoms. use line.
- 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.)
duet rule
H forms stable molecules where it shares only 2 electrons
octet rule
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
configuration tips
noble gas valences from top to bottom:
1s2
2s22p6
3s23p6
4s24p6
5s25p6
6s26p6
evidence of a chemical reaction
color change
formation of a precipitate (solid)
formation of a gas (bubbles)
heat is produced (exothermic) or heat is absorbed (endothermic)
chemical equation
reactants → products
conservation of mass
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
physical states
g - gas
l - liquid
s - solid
aq - aqueous
what causes reactions?
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
precipitation
formation of a solid
solid formed = precipitate
reaction = precipitation reaction
predicting precipitates
insoluble
solid
strong electrolyte
a substance that completely breaks apart into ions when dissolved in water
resulting solution readily conducts an electric current
Ba(NO3)2 and K2CrO4
soluble solid
readily dissolves in water
insoluble solid
slightly soluble solid
only a small amount of the solid dissolves in water
ionic compound
all salts
when ionic compounds dissolve, the resulting solution contains ions
predicting equations
- exchange anions & cations
- balance charges
- balance equation
- use solubility rule to find precipitates
combination
synthesis reaction
two reactants combine to form a single product. the reactants may be elements or compounds
Zn(s) + I2(s) → ZnI2(s)
decomposition
one reactant, a compound, breaks down to give 2+ products
2H2O2(aq) → 2H2O(l) + O2(g)
single replacement
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)
double replacement
2 ionic compounds exchange ions to form new compounds
NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)
combustion
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)
molecular equation
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)
complete ionic equation
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)
spectator ions
ions that don’t participate directly in a reaction
net ionic equation
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]
balancing tip
if 2 elements being combined have the same subscript, u can take it away and change it to a coefficient
atomic mass
molecular weight (MV)
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
amu
atomic mass unit
= 1/12th of the mass of a 12C atom = 1/661 x 10-24 gram
mole
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
formula mass
formula weight FW
total mass for all atoms in a compound
molar mass
the mass (in grams) of 1 mol of the compound and is the sum of the masses of the component atoms
calculating mass using amu
Calculate the mass, in amu, of a sample of aluminum that contains 75 atoms.
- 1 Al atom = 26.98 amu
- 75 Al atoms x (26.98amu/Al atom) = 2024amu
calculating the number of atoms from the mass
Calculate the # of sodium atoms present in a sample that has a mass of 1172.49 amu
- 1 Na atom = 22.99 amu
- 1172.49amu x (1 Na/22.99amu) = 51.00 Na atoms
calculating moles and number of atoms
Compute both the # of moles of atoms and the # of atoms in a 10.0-g sample of aluminum.
- 1 mol Al = 26.98 g Al
- 10.0g Al x (1mol Al/26.98g Al) = 0.371 mol Al
- 6.022 x 1023 Al atoms = 1 mol Al atoms
- 0.371 mol Al x (6.022 x 1023 Al atoms/1 mol Al) = 2.23 x 1023 Al atoms
calculating the number of atoms
How many silicon(Si) atoms are present in 5.68mg? The average atomic mass for Si is 28.09 amu
- 1g = 1000mg
- 5.68mg Si x (1g Si/1000mg Si) = 5.68 x 10-3 g Si
- 1 mol Si atoms = 28.09 g Si
- 5.68 x 10-3 g Si x (1 mol Si/28.09 g Si) = 2.02 x 10-4 mol Si
- 1 mol = 6.022 x 1023
- 2.02 x 10-4 mol Si x (6.022 x 1023 atoms/1 mol Si) = 1.22 x 1020 Si atoms
acid
proton (H+) donor
base
a proton acceptor
acids and bases
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
conjugate acid
formed when the proton is transferred to the base
conjugate acid-base pair
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)
more acids and bases
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
which are acid-base pairs?
- HF, F-→ yes. lost H+
- NH4+, NH3 → yes. lost H+
- HCl, H2O → no
1 & 2 are conjugate acid-base pairs cuz the 2 species differ by one H+
reverse reaction
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
strong acid
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
weak acid
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.
diprotic acid
contributes 2 H+ ions when dissolved in water
Ex: sulfuric acid, H2SO4
amphoteric substance
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.
concentration
[] 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
ion-product constant
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
possible acid situations
in a solution, there are 3 possible situations
- adding an acid to water, increasing [H+]
[H+] > [OH-] → acidic solution
- adding a base to water, increasing [OH-]
[H+] < [OH-] → basic solution
- having equal amounts of acid and base
[H+] = [OH-] → neutral solution
pH
a convenient way to express the acidity of a solution
pH = -log[H+]
1-6 = acidic
7 = neutral
8-14 = basic
