Transcript Ch 1 and Ch 2 PPT with answers

CH 1 CHEMICAL
FOUNDATIONS
AP Chemistry
2014-2015
1.1 CHEMISTRY: AN OVERVIEW
Matter is anything that takes up space and
exhibits inertia.
 Composed of only ~100 types of atoms (ex. water is
made of hydrogen and oxygen; running an electric
current through it separates it into its constituent
elements)
Chemistry is the study of matter and energy—
and more importantly, the changes between
them
1.2 THE SCIENTIFIC METHOD
 Observations
 Measurement = quantitative observation
 No number involved = qualitative observation
 A hypothesis is a possible explanation for an
observation
 Tested by carrying out an experiment
 Repetition of experiments is key
 A theory comes into existence when hypotheses are
assembled in an attempt to explain “why” the “what”
happened
 We use many models to explain natural phenomena
 A scientific law is a summary of observed behavior
 Law of conservation of mass, law of conservation of energy, etc.
1.3 UNITS OF MEASUREMENT
A quantitative observation (measurement)
consists of two parts: a number and a unit
Units
 English (U.S., some of Africa)
 Metric (everyone else, pretty much)
 SI (Le Système International)—developed in 1960 to
improve communication between scientists around
the world; based on/derived from metric system
UNITS CONTINUED
 Volume (derived from length)
 1 dm 3 = 1 L = 1,000 cm 3 = 1,000 mL
 1 cm 3 = 1 mL = 1 g of water at 4°C
 Mass vs. Weight
 Mass (g or kg)—a measure of the resistance of an object to a change
in its state of motion (a measure of inertia); the quantity of matter in
an object
 Weight (Newtons)—the response of mass to gravity
 Gravity varies with altitude; higher at low altitude, lower at high altitude
 Every object has a gravitational field proportional to its mass
 Precision and accuracy
 Two types of error
 See Exercise 1 .2
1.4 SIGNIFICANT FIGURES AND
CALCULATIONS
Rules
 Nonzero digits are significant
 A zero is significant if it is
 Terminating and right of the decimal (must be BOTH)
 “Sandwiched” between significant figures
 Exact or counting numbers have an infinite amount of
significant figures as do fundamental constants
 See Exercise 1.3
RULES FOR CALCULATING WITH SIG FIGS
 Multiplication and division: the term with the least
number of sig figs determines the number of sig figs
in the answer
 Ex. 4.56 x 1.4 = 6.38  6.4
 Addition and subtraction: the term with the least
number of decimal places determines the number of
sig figs in the answer
 Ex. 12.11 + 18.0 + 1.013 = 31.123  31.1
 pH calculations: the number of sig figs in the least
accurate measurements determines the number of
decimal plces on the reported pH
 Round at the end of all calculations
1.6 DIMENSIONAL ANALYSIS
 Consider a pin measuring 2.85 cm in length. What is its
length in inches?
2.54 cm = 1 inch
To convert, multiply your quantity by a conversion factor that
“cancels” the undesired unit and puts the desired unit in its
place.
2.85 cm x (1 inch/2.54 cm) = 1 .12 inches
EXERCISE 1.5 UNIT CONVERSIONS I
EXERCISE 1.6 UNIT CONVERSIONS II
EXERCISE 1.7 UNIT CONVERSIONS III
EXERCISE 1.8 AND 1.9
1.7 TEMPERATURE
 Three scales
 Fahrenheit
 Kelvin
 Celsius
T F = T C x (9°F/5°C) + 32°F
T K = T C + 273.15 K
T C = T K - 273.15 K
 See Exercise 1.10 Temperature Conversions
1.8 DENSIT Y
Density = mass/volume
See Exercise 1.13 Determining Density
1.9 CLASSIFICATION OF MATTER
Solid, liquid, gas
 Fluids = liquids and gases
 Vapor: the gas phase of a substance that is normally
a solid or liquid at room temperature; for example,
we say “water vapor” but we don’t say “oxygen vapor”
CLASSIFICATION CONTINUED
 Mixtures: can be physically separated
 Homogeneous
 Heterogeneous
 Means of physical separation include filtering, fractional
crystallization, distillation, chromatography
 Pure substances: compounds and elements
 Compounds can be separated into elements by chemical
means (ex. electrolysis)
 Elements can be broken down into atoms, which can be
broken down into the nucleus and electron cloud, which
can be broken down into protons, neutrons, and electrons,
which can be broken down into quarks and leptons
CH 2 ATOMS,
MOLECULES, AND IONS
AP Chemistry
2014-2015
2.1 CONTAINS HISTORICAL
INFORMATION.
 Read it if you would like.
2.2 FUNDAMENTAL CHEMICAL LAWS
 Antoine Lavoisier was the fir st chemist to insist on quantitative
experimentation. He got guillotined, but not for that reason.
 The law of conser vation of mass : matter is neither created nor
destroyed.
 The law of definite propor tions : a given compound always contains
exactly the same propor tions of elements by mass.
 The law of multiple proportions : when two
Mass of Oxygen that
elements combine to form a series of
combines with 1 gram of C
compounds, the ratios of the masses of the
Compound I
1.33 g
second element that combine with one gram
2.66 g
of the fir st element can always be reduced to Compound II
small whole number s. You can see an
example of this on the right. The likely formulas for these
compounds would be CO and CO 2.
EXERCISE 2.1 ILLUSTRATING THE LAW OF
MULTIPLE PROPORTIONS
The following data were collected for several compounds of
nitrogen and oxygen:
Mass of Nitrogen That Combines With 1 g of Oxygen
Compound A 1 .750 g
Compound B 0.8750 g
Compound C 0.4375 g
Show how these data illustrate the law of multiple proportions.
2.3 DALTON’S ATOMIC THEORY
 Dalton’s Theory (partially correct, partially not)
 All matter is made of atoms. These indivisible and indestructible
objects are the ultimate chemical particles.
 All the atoms of a given element are identical, in both weight and
chemical properties. However, atoms of different elements have
different weights and different chemical properties.
 Compounds are formed by the combination of different atoms in the
ratio of small whole numbers.
 A chemical reaction involves only the combination, separation, or
rearrangement of atoms; atoms are neither created nor destroyed in
the course of ordinary chemical reactions
 Two modifications were made when subatomic particles and isotopes
were discovered.
Avogadro’s Hypothesis
At the same temperature and pressure, equal
volumes of different gases contain the same
number of particles .
2.4 EARLY EXPERIMENTS TO
CHARACTERIZE THE ATOM
 The electron
 J.J Thomson found that when high voltage was applied to an
evacuated type, a “ray” he called a cathode ray was produced.
The ray was produced at the electrode (also called the cathode)
and was repelled by the negative pole of an applied electric
field. He postulated that the ray was a stream of negative
particles (now called electrons). He then measured the
deflection of beams of electrons to determine the charge -tomass ratio. Thomson discovered that he could repeat this
deflection and calculations using different metal electrodes,
showing that all metals contain electrons and all atoms
contain electrons. He also deduced that since atoms were
neutral, there must be a positive charge within the atom,
giving rise to the “plum pudding” model.
MILLIKAN’S OIL DROP EXPERIMENT
 Next up, Robert Millikan
sprayed charged oil drops into
a chamber. He halted their
fall (due to gravity) by
adjusting the voltage across
two charged plates. He used
the stop-drop voltage and
Thomson’s charge-mass ratio
to determine the charge on
one drop of oil, which was a
whole number multiple of the
electron charge.
 The mass of an electron is
9.11 x 10 -31 kg.
RADIOACTIVIT Y
 Henry Becquerel famously (and accidentally) discovered
radiation when he left a uranium ore in a closed drawer with a
photographic plate. When he realized that the plate had been
exposed, he realized that a form of radiation other than light
had penetrated it. The uranium, of course, was the culprit.
RADIOACTIVIT Y CONTINUED
 Three types of radioactive emission
 Alpha (particles): helium nuclei, relatively massive and slow, poorly
penetrating, somewhat dangerous
 Beta (particles): electrons, relatively light and fast, moderately
penetrating, a little more dangerous
 Gamma (rays): just energy, most penetrating, most dangerous
 These are not the only kinds of radioactive emission. We will discuss
more in the spring.
THE NUCLEAR ATOM
 Rutherford’s famous gold foil experiment proved that
a positively-charged and somewhat bulky nucleus
could be found in the center of an atom. He also
found that atoms are mostly empty space.
2.5 THE MODERN VIEW OF ATOMIC
STRUCTURE (AN INTRODUCTION)
 Elements
 All matter composed of only one type of atom is an element.
92 elements are naturally-occurring; the rest are manmade.
 Atoms
 The atom is the smallest particle of an element that retains
the chemical properties of that element. It consists of a bulky,
dense nucleus (protons and neutrons) and electrons
shells/clouds (which of course contain electrons).
ATOMS AND ISOTOPES
 We can find a few pieces of information about each
element using isotope notation.
 Mass number = #protons + #neutrons for specific isotopes of
an element
 Actual mass is not an integral number! mass defect --causes this
and is related to the energy binding the particles of the nucleus
together
 Atomic number = #protons = #electrons in a neutral atom =
identity of the element
EXERCISE 2.2 WRITING THE SYMBOLS
FOR ATOMS
Write the symbol for the atom that has an
atomic number of 9 and a mass number of 19.
How many electrons and how many neutrons
does this atom have?
ISOTOPES
 Isotopes are atoms that have the same number of
protons (and therefore are the same element) but
different numbers of neutrons (and therefore
different masses).
 Most elements have at least two stable isotopes. Exceptions
include Al, F, P.
 Hydrogen isotopes are important because they have special
names.
 0 neutrons = hydrogen
 1 neutron = deuterium
 2 neutrons = tritium
2.6 MOLECULES AND IONS
 Electrons are responsible for bonding and
chemical reactivity.
 Chemical bonds—forces that hold atoms together
 Covalent bonds—atoms share electrons and make
molecules [independent units]; H 2 , CO 2 , H 2 O, NH 3 , O 2 , CH 4
to name a few.
 Molecule--smallest unit of a compound that retains the
chem. characteristics of the compound; characteristics of
the constituent elements are lost.
 Molecular formula--uses symbols and subscripts to
represent the composition of the molecule. (Strictest
sense--covalently bonded)
MOLECULES AND IONS CONTINUED
 Structural formula—bonds are shown by lines [representing shared e pairs]; do not always indicate shape
 Ions--formed when electrons are lost or gained in ordinary chem.
reactions; dramatically affect size of atom
 Cations--(+) ions; often metals since metals lose electrons to become
+ charged
 Anions--(-) ions; often nonmetals since nonmetals gain electrons to
become - charged
 Polyatomic ions--units of atoms behaving as one entity --MEMORIZE
formula and charge!
 Ionic solids—Electrostatic forces hold ions together. Strong ions held
close together solids.
2.7 AN INTRODUCTION TO THE PERIODIC
TABLE
 Metals—malleable,
ductile & have luster;
most of the elements are
metals—exist as cations
in a “sea of electrons”
which accounts for their
excellent conductive
properties; form oxides
[tarnish] readily and
form POSITIVE ions
[cations]. Why must
some have such goofy
symbols?
PERIODIC TABLE CONTINUED
 Groups or families--vertical columns; have similar
physical and chemical properties (based on similar
electron configurations!!)
 Group A—Representative elements
 Group B--transition elements; all metals; have numerous
oxidation/valence states
 Periods --horizonal rows; progress from metals to
metalloids [either side of the black “stair step” line
that separates metals from nonmetals] to nonmetals
MEMORIZE
ALKALI
METALS—1A
ALKALINE
EARTH
METALS—2A
HALOGENS—7A
NOBLE (RARE)
GASES—8A
2.8 NAMING SIMPLE COMPOUNDS
Binary Ionic Compounds (Type I and Type II)
 In general—consist of a metal cation and a nonmetal
anion.
 The cation is written first. The charges from the cation and
anion must cancel; we use subscripts to make this happen.
 The names of ionic compounds do not contain prefixes such
as mono- or di- unless that is part of the name of a
polyatomic ion in the compound.
 Monatomic ions end in –ide. Ex. NaF is sodium fluoride.
T YPE I BINARY IONIC COMPOUNDS
 Type I contain non-transition metals, which have only
one charge when they are cations.
 Group 1A = +1, Group 2A = +2, Aluminum = +3
 Zinc, silver, and cadmium also fit into this category; silver ions
always have a +1 charge, while zinc and cadmium ions always
have a +2 charge.
 Writing the name of a Type I Binary Ionic compound is simple.
Ex. MgCl 2 is magnesium chloride. The formulas are also
simple, but you have to swap-and-drop to get the correct
formula. Ex. sodium oxide is Na 2 O, and calcium nitride is
Ca 3 N 2 .
T YPE II BINARY IONIC COMPOUNDS
 Type II contain transition metals, as well as a few
others such as lead, tin, and mercury.
 These ions have variable charges which are reflected in
the formula using roman numerals. For example, FeCl 3
would be iron (III) chloride and SnO 2 would be tin (IV)
oxide. Conversely, lead (II) chloride would be PbCl 2 .
 Some of them are real weirdoes. For example, the
mercury (II) ion is Hg 2+ which makes sense, but the
mercury (I) ion is Hg 2 2+ .
EXERCISE 2.3 NAMING T YPE I BINARY
COMPOUNDS
Name each binary compound:
a.CsF
b.AlCl 3
c.LiH
EXERCISE 2.4 NAMING T YPE II BINARY
COMPOUNDS
Give the systematic name of each of the
following compounds.
a. CuCl
b. HgO
c. Fe 2O 3
d. MnO 2
e. PbCl 2
EXERCISE 2.5 NAMING BINARY
COMPOUNDS
Give the systematic name of each of
the following compounds.
a. CoBr 2
b. CaCl 2
c. Al 2O 3
d. CrCl 3
IONIC COMPOUNDS WITH POLYATOMIC
IONS
Same as the other ionic names/formulas
we’ve seen, but you need to look out for
polyatomic ions. I’ve given you a sheet
of them, but you will not be given a list
for the AP Exam.
EXERCISE 2.6 NAMING COMPOUNDS
CONTAINING POLYATOMIC IONS
Give the systematic name of each of the
following compounds.
a. Na 2SO 4
b. KH 2PO 4
c. Fe(NO 3) 3
d. Mn(OH) 2
e. Na 2SO 3
EXERCISE 2.6 CONTINUED
a. Na 2CO 3
b. NaHCO 3
c. CsClO 4
d. NaOCl
e. Na 2SeO 4
f. KBrO 3
BINARY COVALENT COMPOUNDS
Consist of two nonmetals bonded
together
Use prefixes: mono, di, tri, tetra, penta,
hexa, hepta, octa, nona, deca
Don’t forget the –ide ending
EXERCISE 2.7 NAMING T YPE III BINARY
COMPOUNDS
Name each of the following compounds.
a. PCl 5
b. PCl 3
c. SF 6
d. SO 3
e. SO 2
f. CO 2
ACIDS
Hydrogen is listed first in the formula;
the anion is listed second
-ide →hydro [negative ion root]ic ACID
-ate →-ic ACID
-ite → -ous ACID
EXERCISE 2.8 NAMING ACIDS
Give the systematic name for each of the
following acids.
a. H 2 SO 4
b. HClO 3
c. HNO 3
d. H 3 PO 4
e. HCl
f. H 2 CO 3
g. H 2 SeO 3
h. HBrO 2
EXERCISE 2.9 WRITING ACID FORMULAS
Give the formula for each of the following
acids.
a. Hydrobromic acid
b. Perchloric acid
c. Sulfurous acid
d. Acetic acid
e. Iodic acid
f. Dichromic acid
ANNOYING THINGS THAT PEOPLE CAN’T
LET GO OF
Water (easy)
Ammonia NH 3
Hydrazine N 2H 4
Phosphine PH 3
Nitric oxide NO
Nitrous oxide (laughing gas) N 2O
EXERCISE 2.10 NAMING VARIOUS T YPES
OF COMPOUNDS
Give the systematic name for each of the
following compounds.
a. P 4O 10
b. Nb 2O5
c. Li 2O 2
d. Ti(NO 3) 4
EXERCISE 2.11 WRITING COMPOUND
FORMULAS FROM NAMES
Given the following systematic names,
write the formula for each compound.
a. Vanadium(V) fluoride
b. Dioxygen difluoride
c. Rubidium peroxide
d. Gallium oxide