ch 1 chemical foundations
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Ch 1 Chemical Foundations. AP Chemistry 2014-2015. 1.1 Chemistry: An Overview. Matter is anything that takes up space and exhibits inertia. - PowerPoint PPT PresentationTRANSCRIPT
AP Chemistry 2014-2015
CH 1 CHEMICAL FOUNDATIONS
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
1.1 CHEMISTRY: AN OVERVIEW
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.
1.2 THE SCIENTIFIC METHOD
A quantitative observation (measurement) consists of two parts: a number and a unit
UnitsEnglish (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
1.3 UNITS OF MEASUREMENT
Volume (derived from length) 1 dm3 = 1 L = 1,000 cm3 = 1,000 mL 1 cm3 = 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 accuracyTwo types of errorSee Exercise 1.2
UNITS CONTINUED
RulesNonzero digits are significantA 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
1.4 SIGNIFICANT FIGURES AND CALCULATIONS
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
RULES FOR CALCULATING WITH SIG FIGS
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
1.6 DIMENSIONAL ANALYSIS
EXERCISE 1.5 UNIT CONVERSIONS I
EXERCISE 1.6 UNIT CONVERSIONS II
EXERCISE 1.7 UNIT CONVERSIONS III
EXERCISE 1.8 AND 1.9
Three scales Fahrenheit TF = TC x (9°F/5°C) + 32°F
Kelvin TK = TC + 273.15 K
Celsius TC = TK - 273.15 K
See Exercise 1.10 Temperature Conversions
1.7 TEMPERATURE
Density = mass/volume
See Exercise 1.13 Determining Density
1.8 DENSITY
Solid, liquid, gasFluids = liquids and gasesVapor: 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”
1.9 CLASSIFICATION OF MATTER
Mixtures: can be physically separated Homogeneous Heterogeneous Means of physical separation include filtering,
fractional crystallization, distillation, chromatographyPure 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
CLASSIFICATION CONTINUED
AP Chemistry 2014-2015
CH 2 ATOMS, MOLECULES, AND
IONS
Read it if you would like.
2.1 CONTAINS HISTORICAL INFORMATION.
Antoine Lavoisier was the fi rst chemist to insist on quantitative experimentation. He got guillotined, but not for that reason.
The law of conservation of mass : matter is neither created nor destroyed.
The law of defi nite proportions : a given compound always contains exactly the same proportions of elements by mass.
The law of multiple proportions : when two elements combine to form a series of compounds, the ratios of the masses of the second element that combine with one gram of the fi rst element can always be reduced to small whole numbers. You can see an example of this on the right. The likely formulas for these compounds would be CO and CO 2 .
2.2 FUNDAMENTAL CHEMICAL LAWS
Mass of Oxygen that combines with 1
gram of CCompoun
d I1.33 g
Compound II
2.66 g
The following data were collected for several compounds of nitrogen and oxygen: Mass of Nitrogen That Combines With 1 g of OxygenCompound A 1.750 g Compound B 0.8750 g Compound C 0.4375 g Show how these data illustrate the law of multiple proportions.
EXERCISE 2.1 ILLUSTRATING THE LAW OF MULTIPLE
PROPORTIONS
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.
2.3 DALTON’S ATOMIC THEORY
Avogadro’s HypothesisAt the same temperature and pressure, equal volumes of different gases contain the same number of particles .
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-to-mass 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.
2.4 EARLY EXPERIMENTS TO CHARACTERIZE THE ATOM
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.
MILLIKAN’S OIL DROP EXPERIMENT
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.
RADIOACTIVITY
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.
RADIOACTIVITY CONTINUED
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.
THE NUCLEAR ATOM
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).
2.5 THE MODERN VIEW OF ATOMIC STRUCTURE (AN INTRODUCTION)
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
ATOMS AND ISOTOPES
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?
EXERCISE 2.2 WRITING THE SYMBOLS FOR ATOMS
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
ISOTOPES
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]; H2, CO2, H2O, NH3, O2, CH4 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)
2.6 MOLECULES AND IONS
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.
MOLECULES AND IONS CONTINUED
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?
2.7 AN INTRODUCTION TO THE PERIODIC TABLE
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
PERIODIC TABLE CONTINUED
ALKALI METALS—1A
ALKALINE EARTH METALS—2A
HALOGENS—7A
NOBLE (RARE) GASES—8A
MEMORIZE
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.
2.8 NAMING SIMPLE 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. MgCl2 is magnesium chloride. The formulas are also simple, but you have to swap-and-drop to get the correct formula. Ex. sodium oxide is Na2O, and calcium nitride is Ca3N2.
TYPE I 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, FeCl3 would be iron (III) chloride and SnO2 would be tin (IV) oxide. Conversely, lead (II) chloride would be PbCl2.
Some of them are real weirdoes. For example, the mercury (II) ion is Hg2+ which makes sense, but the mercury (I) ion is Hg2
2+.
TYPE II BINARY IONIC COMPOUNDS
Name each binary compound: a.CsF b.AlCl3
c.LiH
EXERCISE 2.3 NAMING TYPE I BINARY COMPOUNDS
Give the systematic name of each of the following compounds. a. CuCl b. HgO c. Fe2O3
d. MnO2
e. PbCl2
EXERCISE 2.4 NAMING TYPE II BINARY COMPOUNDS
Give the systematic name of each of the following compounds. a.CoBr2
b.CaCl2
c.Al2O3
d.CrCl3
EXERCISE 2.5 NAMING BINARY COMPOUNDS
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.
IONIC COMPOUNDS WITH POLYATOMIC IONS
Give the systematic name of each of the following compounds. a. Na2SO4
b. KH2PO4
c. Fe(NO3)3
d. Mn(OH)2
e. Na2SO3
EXERCISE 2.6 NAMING COMPOUNDS CONTAINING
POLYATOMIC IONS
a.Na2CO3
b.NaHCO3
c.CsClO4
d.NaOCl e.Na2SeO4
f. KBrO3
EXERCISE 2.6 CONTINUED
Consist of two nonmetals bonded together
Use prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca
Don’t forget the –ide ending
BINARY COVALENT COMPOUNDS
Name each of the following compounds. a. PCl5
b. PCl3
c. SF6
d. SO3
e. SO2
f. CO2
EXERCISE 2.7 NAMING TYPE III BINARY COMPOUNDS
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
ACIDS
Give the systematic name for each of the following acids.a. H2SO4
b. HClO3
c. HNO3
d. H3PO4
e. HClf. H2CO3
g. H2SeO3
h. HBrO2
EXERCISE 2.8 NAMING ACIDS
Give the formula for each of the following acids.a. Hydrobromic acidb. Perchloric acidc. Sulfurous acidd. Acetic acide. Iodic acidf. Dichromic acid
EXERCISE 2.9 WRITING ACID FORMULAS
Water (easy)Ammonia NH3
Hydrazine N2H4
Phosphine PH3
Nitric oxide NONitrous oxide (laughing gas) N2O
ANNOYING THINGS THAT PEOPLE CAN’T LET GO OF
Give the systematic name for each of the following compounds. a. P4O10
b. Nb2O5
c. Li2O2
d. Ti(NO3)4
EXERCISE 2.10 NAMING VARIOUS TYPES OF COMPOUNDS
Given the following systematic names, write the formula for each compound. a. Vanadium(V) fluorideb. Dioxygen difluoridec. Rubidium peroxided. Gallium oxide
EXERCISE 2.11 WRITING COMPOUND FORMULAS FROM
NAMES