Summer Study - AP Chemistry - Chemistry - 2011-09-09

1.4—Classifications of Matter

Matter is anything that occupies space.

Substances, mixtures, elements, compounds, atoms, and molecules are all forms of matter.

A substance is a form of matter that has a definite (constant) composition and distinct properties.

A mixture is a combination of two or more substances in which the substances retain the distinct identities.

In a homogenous mixture the composition of the mixture is the same throughout.

In a heterogeneous mixture the composition is not uniform.

An element is a substance that cannot be separated into simpler substances by chemical means.

A compound is a substance composed of atoms of two or more elements chemically united in fixed proportions.

 

1.6—Physical and Chemical Properties of Matter

A physical property can be measured and observed without changing the composition or identity of a substance.

            The melting point of a substance is a physical property.

The measured value of an extensive property depends on how much matter is being considered.

Mass, which is the quantity of matter in a given sample of a substance, is an extensive property.

            Values of the same extensive property can be added together (?).

            Volume is also an extensive property.

The measured value of an intensive property does not depend on how much matter is being considered.

Density, defined as the mass of an object divided by its volume, is an intensive property.

Temperature is also an intensive property.

Unlike mass, length, and volume, temperature and other intensive properties are not additive.

 

1.7—Measurement

Macroscopic properties can be determined directly.

Microscopic properties, on the atomic of molecular scale, must be determined by an indirect method.

Technically, weight is the force that gravity exerts on an object.

A liter (L) is the volume occupied by one cubic decimeter.

1 liter of volume is equal to 1000 milliliters (mL) or 1000 cm3.

1 mL = 1 cm3

density = mass/volume or d = m/V

            As m increases, so does V.

            Density usually decreases with temperature.

1 g/cm3 = 1g/mL = 1000 kg/m3

1g/L = 0.001 g/mL

The Kelvin scale does not have the degree sign. Also, temperatures expressed in kelvins can never be negative because theoretically, 0 K is the lowest temperature that can ever be attained.

To convert degrees Fahrenheit to degrees Celsius:

            ?°C = (°F -32°F) x 5°C/9°F

To convert degrees Celsius to degrees Fahrenheit:

            ? °F = 9°F/5°C x (°C) + 32°F

I Kelvin = 1°Celsius and absolute zero on the Kelvin scale is equal to -273.15°C.

To convert Celsius to Kelvin:

? K = (°C + 273.15°C) 1 K/1°C

 

1.8—Handling Numbers

Accuracy tells us how close a measurement is to the true value of the quantity that was measured.

Precision refers to how closely two or more measurements of the same quantity agree with one another.

 

 

Chapter 2

 

2.1—The Atomic Theory

Democritus- believed that there were small, indivisible particles, which he named atomos.

In 1808 John Dalton formed a precise definition of the building blocks of matter that we call atoms.

Dalton’s Atomic Theory

  1. Elements are composed of small particles called atoms.
  2. All atoms of a given element are identical, having the same size, mass, and chemical properties. The atoms of one element are different from the atoms of all other elements.
  3. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer of a simple fraction.
  4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction.

Joseph Prout’s law of definite proportions states that different samples of the same compound always contain its constituent elements in the same proportion by mass.

The law of multiple proportions states that if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers. (This is supported by Dalton’s third hypothesis).

Dalton’s 4th hypothesis is the same as the law of conservation of mass, which states that matter can be neither created nor destroyed.

 

2.2—The Structure of an Atom

An atom is the basic unit of an element that can enter into chemical combination.

Scientists later found out that atoms possessed an internal structure made up of even smaller particles, which are known as subatomic particles. This led to the discovery of protons, neutrons, and electrons.

The Electron

In the 1890’s, many scientists became caught up in the study of radiation, which is the emission and transmission of energy through space in the form of waves.

The cathode ray tube was used to investigate further into the structure of the atom.

English Physicist J.J. Thompson, used a cathode ray tube to determine th ration of electric charge to the mass of an individual electron. The number he came up with was -1.776 x 108 C/g (where C stands for coulonb).

From 1908-1917 R.A. Millikan carried out a series of experiments and succeeded in measuring the charge of an electron to be -1.6602 x 10-19 C.

From that, Millikan determined the mass of an electron to me 9.10 x 10-28 g.

Radioactivity

In 1895, the German physicist Wilhelm Röntgen noticed that cathode rays caused glass and metals to emit very unusual rays. He called them X rays because of their unknown nature.

Marie Curie, one of Antoine Becquerel’s students (he was a physics professor in Paris at the time) suggested the name radioactivity to describe the spontaneous emission of particles and/or radiation.

3 types of rays are produced by the decay, or breakdown, or radioactive substances such as uranium. 2 of the 3 are deflected by oppositely charged metal plates.

            Alpha (a) rays consist of positively charged particles, called a particles, and therefore car deflected by the positively charged plate.

            Beta (b) rays are electrons and are deflected by the negatively charged plate.

The third type of radioactive radiation consists of high-energy rays called gamma (g) rays. Like X rays, gamma rays have no charge and are not affected by and external field.

The Proton and the Nucleus

Thomson proposed that an atom could be thought of as a uniform, positive sphere in which the electrons are embedded into.

In 1910 the New Zealand physicist Ernest Rutherford conducted his gold foil experiment where he used very thin gold foils and other metals as targets for a particles. It surprised him when in some instances the a particles bounced back in the direction from which it had come.

Rutherford later explained his findings in terms of a new model for the atom. He said that most of the atom was empty space. He proposed that the atom’s positive charge was all concentrated in the nucleus, which is a dense central core within the atom.

The positively charged particles in the nucleus are called protons.

In separate experiments, it was found that the proton carries the same quantity of charge as an electron and has a mass of 1.67262 x 10-24 g, which is about 1840 times the mass fo the electron.

Atomic and molecular dimensions are expressed in SI units called picometers (pm) where

            1 pm = 1 x 10-12 m

A typical atomic radius is about 100 pm, whereas the radius of an atomic nucleus is only about 5 x 10-3 pm.

The Neutron

An English physicist named James Chadwick provided the proof that there was a third type of subatomic particle in an atom.

When he bombarded a thin sheet of beryllium with a particles, a very high-energy radiation similar to gamma rays was emitted by the metal. Later experiments show that the rays actually consisted of a third type of subatomic particles, which Cahdwick named neutrons, because they proved to be electrically neutral particles hacing a mass slightly greater and that of protons.

 

2.3—Atomic Number, Mass number, and Isotopes

The atomic number (Z) is the number of protons in the nucleus of each atom of an element.

In a neutral atom the number of protons is equal to the number of electrons so the atomic number also indicates the number of electrons present in the aton.

The mass number (A) is the total number of neutrons and protons present in the nucleus of an atom of an element.

In general, the mass number is given by

            Mass number = number of protons + number of neutrons

                                   = atomic number + number of neutrons

The number of neutrons in an atom is equal to the difference between the ass number and the atomic number, or (A-Z).

Most elements have tow or more isotopes, atoms that have the same atomic number but different mass numbers.

            mass numberà AZ X

            atomic numberà  

2.4—The Periodic Table

The periodic table is a chart in which elements having similar chemical and physical properties are grouped together.

            Elements are arranged by atomic number in horizontal rows called periods and in vertical columns known as groups or families.

Elements can be divided into 3 subcategories.

            A metal is a good conductor of heat and electricity while a nonmetal is usually a poor conductor of heat and electricity. A metalloid has properties that are intermediate between those of metals and nonmetals.

From left to right across any period the physical and chemical properties of the elements change gradually from metallic to nonmetallic.

Group 1A elements (Li, Na, K, Rb, Cs, and Fr) are called alkali metals.

Groups 2A elements (Be, Mg, Ca, Sr, Ba, and Ra) are called alkali earth metals.

Elements in Group 7A (F, Cl, Br, I, and At) are known as halogens.

And elements in Group 8A (He, Ne, Ar, Kr, Xe, Rn) are called noble gases, or rare gases.

 

2.5—Molecules and Ions

Only 6 noble gases occur in nature as single atoms, so they are called monatomic gases.

A molecule is an aggregate of at least two atoms in a definite arrangement held together by chemical forces (also called chemical bonds). Molecules are electrically neutral.

The hydrogen molecule is called a diatomic molecule because it contains only two atoms.

Molecules containing more than two atoms are called polyatomic molecules.

An ion is an atom or a group of atoms that has a net positive or negative charge.

A cation, is an ion with a positive charge.

An anion is an ion whose net charge is negative due to an increase in the number of electrons.

An ionic compound is formed from cations and anions.

Monatomic ions contain only one atom. Some examples are Mg2+, Fe3+, S2-, and N3-.

Usually, metals tend to form cations and nonmetals tend to form anions.

Polyatomic ions, such as OH- are ions containing more than one atom.

 

2.6—Chemical Formulas

Chemical formulas are used to express the composition of molecules and ionic compounds in terms of chemical symbols.

A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance.

An allotrope is one of two or more distinct forms of an element. (Fore example oxygen (O2) and ozone (O3) are allotropes of oxygen).

The structural formula shows how atoms are bonded to one another in a molecule.

The empirical formula tells us which elements are present and the simplest whole-number ratio of their atoms, but not necessarily the actual number of atoms in a given molecule.

Empirical formulas are the simplest chemical formulas.

Molecular formulas are the true formulas of molecules.

In the charges on the cation and anion are numerically different, we apply the following rule to make the formula electrically neutral: The subscript of the cation is numerically equal to the charge on the anion, and the subscript of the anion is numerically equal to the charge on the cation.

 

2.7—Naming Compounds

Organic compounds contain carbon, usually in combination with elements such as hydrogen, oxygen, nitrogen, and sulfur.

All other compounds are classified as inorganic compounds.

 

Ionic Compounds

Metal cations take their names from the elements. For example,

            Na            sodium                        Na+            sodium ion (or sodium cation)

            K            potassium            K+            potassium ion (or potassium cation)

            Mg            magnesium            Mg2+            magnesium ion (or magnesium cation)

            Al            aluminum            Al3+            aluminum ion (or aluminum cation)

Many ionic compounds are binary compounds, or compounds formed from just two elements.

The compounds LiOH and KCN are named litium hydroxide and potassium cyanide, resectfully. These and a number of other such ionic substances are called ternary compounds, meaning compounds consisting of three elements.

It has become increasingly common to designate different cations with Roman numerals. This is called the Stock system.

 

Molecular Compounds

It is quite common for one pair of elements to form several different compounds so to avoid confusion Greek prefixes are used. For example:

            CO            carbon monoxide

            CO2            carbon dioxide

            SO2            sulfur dioxide

            SO3            sulfur trioxide

            NO2            nitrogen dioxide

            N2O4            dinitrogen tetroxide

 

The prefixes are:

Prefix

Meaning

Mono-

1

Di-

2

Tri-

3

Tetra-

4

Penta-

5

Hexa-

6

Hepta-

7

Octa-

8

Nona-

9

Deca-

10

 

The prefix “mono-” may be omitted for the first element.

For oxides, the ending “a” in the prefix is sometimes omitted. For example, N2O4 may be called dinitrogen tetroxide rather than dinitrogen tetraoxide.

Exceptions to the use of Greek prefixes are molecular compounds containing hydrogen.

            B2H6            diborane

            CH4                  methane

            SiH4            silane

            NH3            ammonia

            PH3            phosphine

            H2O            water

            H2S            hydrogen sulfide

 

Acids and Bases

An acid is a substance that yields hydrogen ions (H+) when dissolved water in water. (H+ is equivalent to one proton and is often referred to that way).

Formulas for acids contain one or more hydrogen atoms as well as an anionic group.

Anions whose names end in “-ide” form acids with a “hydro-” prefix and an “-ic” ending.

In some cases two different names seem to be assigned t the same chemical formula.

            HCl            hydrogen chloride

            HCl            hydrochloric acid

Oxoacids are acids that contain hydrogen, oxygen, and another element (the central element).

            The formulas of oxoacids are usually written with the H first, followed by the central element and then O.

We use the following five common acids as out references in naming oxoacids:

            H2CO3                        carbonic acid

            HClO3                        chloric acid

            HNO3                        nitric acid

            H3PO4                        phosphoric acid

            H2SO4                        sulfuric acid

Often two or more oxoacids have the same central atom but a different number of O atoms. Starting with our reference oxoacids whose names all end with “-ic”, we use the following rules to mane these compounds.

  1. Addition of one O atom to the “-ic” acid: The acid is “per … -ic” acid. Thus, adding an O atom to HClO3 changes chloric acid to perchloric acid, HClO4.
  2. Removal of one O atom from the “-ic” acid: The acid is called “-ous” acid. Thus nitric acid, HNO3, becomes nitrous acid, HNO2.
  3. Removal of two O atoms from the “-ic” acid: The acid is called “hypo … -ous” acid. Thus, when HBrO3 is converted to HBrO, the acid is called hypobromous acid.

The rules for naming oxoanions, anions of oxoacids, are as follows:

  1. When all the H ions are removed from the “-ic” acid, the anion’s name ends with “-ate.” For example, the anion CO32- derived from H2CO3 is called carbonate.
  2. When all the H ions are removed from the “-ous” acid, the anion’s name ends with “-ite.” Thus, the anion ClO2- derived from HClO2 is called chlorite.
  3. The manes of anions in which one or more but not all the hydrogen ions have been removed must indicate the number of H ions present. For example, consider the anions derived from phosphoric acid:

H3PO4                        phosphoric acid

H2PO4-                        dihydrogen phosphate

HPO42-                        hydrogen phosphate

PO43-                        phosphate

Naming Bases

A base can be described as a substance that yields hydroxide ions (OH--) when dissolved in water. Some examples are

            NaOH                        sodium hydroxide

            KOH                        potassium hydroxide

            Ba(OH)2            barium hydroxide

Ammonia (NH3) is also classified as a common base.

 

Hydrates

Hydrates are compounds that have a specific number of water molecules attached to them.

 

3.1—Atomic Mass

Atomic mass (or atomic weight) is the mass of the atom in atomic mass units (amu).

One atomic mass unit is defined as a mass exactly equal to one-twelfth the mass of one carbon 12 atom.

 

3.2—Avogadro’s Number and the Molar Mass of an Element

In the SI system the mole (mol) is the amount of a substance that contains as many elementary entities (atoms, molecules, or other particles) as there are atoms in exactly 12 g (or 0.012 kg) of the carbon-12 isotope.

            This number is called Avogadro’s number (NA) and the current accepted value is 6.0221415 x 1023.

Molar mass is defined as the mass (in grams or kilograms) of 1 mole of units( such as atoms or molecules) of a substance.

 

3.3—Molecular Mass

The molecular mass (sometimes called molecular weight) is the sum of the atomic masses (in amu) in the molecule.

 

3.5—Percent Composition of Compounds

The percent composition by mass is the percent by mass of each element in a compound.

            Percent composition of an element = (n x molar mass of element) / (molar mass of compound) x 100% where n is the number of moles of the element in 1 mole of the compound.

 

3.7—Chemical Reactions and Chemical equations

A chemical reaction is a process in which a substance (or substances) is changed into one or more new substances.

C chemical equation uses chemical symbols to show what happens during a chemical reaction.

            2H2 + O2 à 2H2O

In this equation we refer to H2 and O2 as reactants, which are the starting materials in a chemical reaction. Water is the product, which is the substance formed as a result of a chemical reaction.

In a chemical equation, the reactants are usually written on the left and the products on the right side of the arrow.

            Reactants à Products

 

3.8—Amounts of Reactants and Products

Stoichiometry is the quantitative study of reactants and products in a chemical reaction.

Whatever the given units for the reactants, we use mole to calculate the amount of product formed in a reaction. This approach is called the mole method, which means simply that the stoichiometric coefficients in a chemical equation cam be interpreted as the number of moles of each substance.

The general approach for solving stoichiometric problems is:

  1. Write a balanced equation for the reaction.
  2. Convert the given amount of the reactant (in grams or other units) to number of moles.
  3. Use the mole ratio from the balanced equation to calculate the number of moles of product formed.
  4. Convert the moles of product to grams (or other units) or product.

 

3.9—Limiting Reagents

When a chemist carries out a reaction, the reactants are usually not present in exact stoichiometric amounts, that is, in the proportions indicated by the balanced equation.

The reactant used up first in a reaction is called the limiting reagent, because the maximium amount of product formed depends on how much of this reactant was originally present.

Excess reagents are the reactants present in quantities greater than necessary to react with the quantity of the limiting reagent.

 

3.10—Reaction Yield

The amount of limiting reagent present at the start of a reaction determines the theoretical yield of the reaction, that is, the amount of product that would result if all the limiting reagent reacted.

            The theoretical yield, then, is the maximum obtainable yield, predicted by the balanced equation.

In practice, the actual yield, or the amount of product actually obtained from a reaction, is almost always less than the theoretical yield.

To determine how efficient a given reaction is, chemist often figure the percent yield, which describes the proportion of the actual yield to the theoretical yield.

            % yield = (actual yield / theoretical yield) x 100%

 

5.4—The Ideal Gas Equation

Boyle’s Law: P1V2 = P2V2(at constant n and T)

Charles’s Law: V1/T1 = V2/T2 (at constant n and P)

Avogadro’s Law: V = kn (at constant P and T)

We can combine all three expressions to form a single master equation for the behavior of gasses:

                        PV = nRT

where R, the proportionality constant, is called the gas constant. This equation is known as the ideal gas equation, which describes the relationship among the four variables P, V, T and n.

An idea gas is a hypothetical gas whose pressure-volume-temperature behavior can be completely accounted for by the ideal gas equation.

The conditions 0°C and 1 atm are called standard temperature and pressure, or STP.

 

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