Chemistry 1A Book Notes

Page 1

Chapter 1

Chemical Foundations

Overview of Chemistry

Matter is composed of various types of atoms.

One substance changes to another by re-organizing the way the atoms are attached to each other.

Scientific method: Observation>Hypothesis>Experiment>Theory (Can be modified, also called a model)

Observations may be qualitative (the sky is blue; water is a liquid) or quantitative (water boils at 1008C; a certain chemistry book weighs 2 kg).

A law summarizes what happens; a theory (model) is an attempt to explain why it happens

Units of Measure

A quantitative observation, or measurement, always consists of two parts: a number and a scale (called a unit).

Si unit= based on metric system, used by scientists.

(1 m)^3= (10 dm)^3= 1000 dm3. A cubic decimeter, that is, (1 dm)3, is commonly called a liter.

1 cm35 1 milliliter (mL) ( 1 L51000 cm351000 mL )

Mass and weight are two different terms.

Uncertainty in Measurement

Measure at bottom of meniscus.

Uncertain digit is the last digit.

The quantity 25 mL means that the volume is between 24 mL and 26 mL, whereas the quantity 25.00 mL means that the volume is between 24.99 mL and 25.01 mL.

Accuracy= closeness to true value

Precision= closeness of results.

Significant Figures

Non-zero #'s are always significant.

Trailing zeros are only significant if it has a decimal.

Leading zeros that proceed nonzero #'s are insignificant.

Zeros between sigfigs are significant.

Exact numbers are usually obtained from counting (3 apples, 2 molecules) and have an infinite # of sigfigs.

For multiplication or division, use smallest number of sigfigs. EX: 1.1*2=2.2 -> 2

For addition or subtraction, answer should be least amount of decimals. EX: 23.445+7.83=31.275 -> 31.28

Always round at the end of the calculation if needed.

Dimensional Analysis

Conversion of one or more unit factor to another using division. Some conversion factors:

1m=1.094yd

2.54cm=1in

1kg=2.205lb

453.6=1lb

1L=1.06qt

1ft^3=28.32L

Temperature

Temp(K) = Temp(C) + 273.15

Temp(C) = Temp(K)-273.15

Temp(F) = Temp(C) *9f/5c+ 32F

Temp(C)= (Temp(F)-32F)(5C/9F)

Density

Density=Mass/Volume

Page 2

Chapter 2

Atoms, Molecules, and Ions

Classification of Matter

Matter= anything that has occupies space and has mass.

Solids: fixed volume and shape

Liquid: fixed volume, unfixed shape

Gas: no fixed volume or shape

Fundamental Chemical Laws

Law of conservation of mass: Mass is neither created nor destroyed in a chemical reaction. ( Antoine Lavoisier )

Law of definite proportion: A given com-pound always contains exactly the same proportion of elements by mass.

Law of multiple proportions: When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 g of the first element can

Dalton's Atomic Theory

All elements are composed of atoms.

All atoms of a given element are identical.

Chemical compounds are formed when atoms combine.

Atoms are not changed in chemical reactions, but the way they are bound together changes

Atom

Protons = +1 charge, in nucleus

Neutrons= 0 charge, in nucleus

Electron= -1 charge, outside

Page 3

Chapter 3

Stoichiometry

Deals with the amounts of substances consumed and/or produced in a chemical reaction.

To relate mass and the number of atoms, the average atomic mass is required.

Mole

The mass of 1 mole of an element = the atomic mass in grams.

The amount of carbon atoms in exactly 12 g of pure 12^C

6.022 * 10^23 units of a substance

Molar mass

Mass (g) of 1 mole of a compound or element

Obtained for a compound by finding the sum of the average masses of its constituent atoms.

Percent composition

The mass percent of each element in a compound.

Mass percent = (mass of element in 1 mole of substance/mass of 1 mole of substance)*100

Empirical formula

The simplest whole-number ratio of the various types of atoms in a compound

Can be obtained from the mass percent of elements in a compound

Molecular formula

For molecular substances: The formula of the constituent molecules

Always an integer multiple of the empirical formula

For ionic substances: The same as the empirical formula

Chemical reactions

Reactants are turned into products.

Atoms are neither created nor destroyed.

All of the atoms present in the reactants must also be present in the products.

Characteristics of a chemical equation

Represents a chemical reaction

Reactants on the left side of the arrow, products on the right side

When balanced, gives the relative numbers of reactant and product molecules or ions

Stoichiometry calculations

Amounts of reactants consumed and products formed can be determined from the balanced chemical equation.

The limiting reactant is the one consumed first, thus limiting the amount of product that can form.

The theoretical yield is the maximum amount that can be produced from a given amount of the limiting reactant.

The actual yield, the amount of product actually obtained, is always less than the theoretical yield.

Percent yield = (actual yield(g)/theoretical yield(g))*100

Page 4

Chapter 4

Types of Chemical Reactions and Solution Stoichiometry

Water is a polar solvent that dissolves many ionic and polar substances.

A substance, the solute, is dissolved in liquid water, the solvent.

Electrolytes

Strong electrolyte: 100% dissociated to produce separate ions; strongly conducts an electric current

Weak electrolyte: Only a small percentage of dissolved molecules produce ions; weakly conducts an electric current

Nonelectrolyte: Dissolved substance produces no ions; does not conduct an electric current

Acids and bases

Arrhenius model

Acid: produces H1

Base: produces OH2

Bronsted-Lowry model

Acid: proton donor

Base: proton acceptor

Strong acid: completely dissociates into separated H1 and anions

Weak acid: dissociates to a slight extent

Molarity

One way to describe solution composition

M=moles/liters

Dilution

Solvent is added to reduce the molarity

Moles of solute after dilution = moles of solute before dilution

M1V1 = M2V2

Types of equations that describe solution reactions

Formula equation: All reactants and products are written as complete formulas

Complete ionic equation: All reactants and products that are strong electrolytes are written as separated ions

Net ionic equation: Only those compounds that undergo a change are written; spectator ions are not included

Solubility rules

Based on experiment observation

Help predict the outcomes of precipitation reactions

Important types of solution reactions

Acid–base reactions: involve a transfer of H1 ions

Precipitation reactions: formation of a solid occurs

Oxidation–reduction reactions: involve electron transfer

Titrations

Measures the volume of a standard solution (titrant) needed to react with a substance in solution

Stoichiometric (equivalence) point: the point at which the required amount of titrant has been added to exactly react with the substance being analyzed

Endpoint: the point at which a chemical indicator changes color

Oxidation–reduction reactions

Oxidation states are assigned using a set of rules to keep track of electron flow.

Oxidation: increase in oxidation state (a loss of electrons)

Reduction: decrease in oxidation state (a gain of electrons)

Oxidizing agent: gains electrons (is reduced)

Reducing agent: loses electrons (is oxidized)

Equations for oxidation–reduction reactions can be balanced by the oxidation states method.

Page 5

Chapter 5

Gases

The state of a gas can be described completely by specifying its pressure (P), volume (V), temperature (T), and the amount (moles) of gas present (n)

Gas laws

Boyle’s law: PV = k

Charles’s law: V = T

Avogadro’s law: V = n

Ideal gas law: PV = nRT

Dalton’s law of partial pressures: Ptotal = P1 + P2 + P3 … where Pn represents the partial pressure of component n in a mixture of gases

Pressure: 1 torr=1mmHg, 1atm=760torr, 1atm=101325Pa

Kinetic molecular theory

Model that accounts for ideal gas behavior

Volume of gas particles is zero

No particle interactions

Particles are in constant motion, colliding with the container walls to produce pressure

The average kinetic energy of the gas particles is directly proportional to the temperature of the gas in kelvins

Gas properties

The particles in any gas sample have a range of velocities

The root mean square (rms) velocity for a gas represents the average of the squares of the particle velocities

Diffusion: the mixing of two or more gases

Effusion: the process in which a gas passes through a small hole into an empty chamber

Real gas behavior

Real gases behave ideally only at high temperatures and low pressures

Understanding how the ideal gas equation must be modified to account for real gas behavior helps us understand how gases behave on a molecular level

Van der Waals found that to describe real gas behavior we must consider particle interactions and particle volumes

Page 6

Chapter 6

Thermochemistry

Energy

The capacity to do work or produce heat

Is conserved (first law of thermodynamics)

Can be converted from one form to another

Is a state function

Potential energy: stored energy

Kinetic energy: energy due to motion

The internal energy for a system is the sum of its potential and kinetic energies

The internal energy of a system can be changed by work and heat: ∆E=q+w

Work

Force applied over a distance

For an expanding/contracting gas

Not a state function

w=-P∆V

Heat

Energy flow due to a temperature difference

Exothermic: energy as heat flows out of a system

Endothermic: energy as heat flows into a system

Not a state function

Measured for chemical reactions by calorimetry

Enthalpy

H=E+PV

Is a state function

Hess’s law: the change in enthalpy in going from a given set of reactants to a given set of products is the same whether the process takes place in one step or a series of steps

Page 7

Chapter 7

Atomic structure and periodicity

Electromagnetic radiation

Characterized by its wavelength λ, frequency v, and speed c=2.9979*10^8 m/s

Can be viewed as a stream of “particles” called photons, each with energy hv, where h is Planck’s constant (6.626*10^-34J*s)

Photoelectric effect

When light strikes a metal surface, electrons are emitted

Analysis of the kinetic energy and numbers of the emitted electrons led Einstein to suggest that electromagnetic radiation can be viewed as a stream of photons

Hydrogen spectrum

The emission spectrum of hydrogen shows discrete wavelengths

Indicates that hydrogen has discrete energy levels

Bohr model of the hydrogen atom

Using the data from the hydrogen spectrum and assuming angular momentum to be quantized, Bohr devised a model in which the electron traveled in circular orbits

Although an important pioneering effort, this model proved to be entirely incorrect

Wave (quantum) mechanical model

An electron is described as a standing wave

The square of the wave function (often called an orbital) gives a probability distribution for the electron position

The exact position of the electron is never known, which is consistent with the Heisenberg uncertainty principle: It is impossible to know accurately both the position and the momentum of a particle simultaneously

Probability maps are used to define orbital shapes

Orbitals are characterized by the quantum numbers n, l, and ml.

Electron spin

Described by the spin quantum number ms, which can have values of土0.5

Pauli exclusion principle: No two electrons in a given atom can have the same set of quantum numbers n, l, ml, and ms

Only two electrons with opposite spins can occupy a given orbital

Periodic table

By populating the orbitals from the wave mechanical model (the aufbau principle), the form of the periodic table can be explained

According to the wave mechanical model, atoms in a given group have the same valence (outer) electron configuration

The trends in properties such as ionization energies and atomic radii can be explained in terms of the concepts of nuclear attraction, electron repulsions, shielding, and penetration

Page 8

Chapter 8

Bonding: general concepts

Chemical bonds

Hold groups of atoms together

Occur when a group of atoms can lower its total energy by aggregating

Types of chemical bonds

Ionic: electrons are transferred to form ions

Covalent: equal sharing of electrons

Polar covalent: unequal electron sharing

Percent ionic character of a bond X-Y: (measured dipole moment of X-Y/calculated dipole moment for X^+Y^-)*100

Ionic bonding

An ion has a different size than its parent atom

An anion is larger than its parent ion

A cation is smaller than its parent atom

Lattice energy: the change in energy when ions are packed together to form an ionic solid

Bond energy

The energy necessary to break a covalent bond

Increases as the number of shared pairs increases

Can be used to estimate the enthalpy change for a chemical reaction

Lewis structures

Show how the valence electron pairs are arranged among the atoms in a molecule or polyatomic ion ❯ Stable molecules usually contain atoms that have their valence orbitals filled

Leads to a duet rule for hydrogen

Leads to an octet rule for second-row elements

The atoms of elements in the third row and beyond can exceed the octet rule

Several equivalent Lewis structures can be drawn for some molecules, a concept called resonance

When several nonequivalent Lewis structures can be drawn for a molecule, formal charge is often used to choose the most appropriate structure(s)

VSEPR model

Based on the idea that electron pairs will be arranged around a central atom in a way that minimizes the electron repulsions

Can be used to predict the geometric structure of most molecules

Page 9

Chapter 9

Covalent bonding: orbitals

Two widely used bonding models

Localized electron model

Molecular orbital model

Localized electron model

Molecule is pictured as a group of atoms sharing electron pairs between atomic orbitals

Hybrid orbitals, which are combinations of the “native” atomic orbitals, are often required to account for the molecular structure

Six electron pairs (octahedral arrangement) require d^2sp^3 orbitals

Five electron pairs (trigonal bipyramidal arrangement) require dsp3 orbitals

Four electron pairs (tetrahedral arrangement) require sp3 orbitals

Three electron pairs (trigonal planar arrangement) require sp2 orbitals

Two electron pairs (linear arrangement) require sp orbitals

Two types of bonds

Sigma σ: electrons are shared in the area centered on a line joining the atoms

Pi π: a shared electron pair occupies the space above and below the line joining the atoms

Molecular orbital model

A molecule is assumed to be a new entity consisting of positively charged nuclei and electrons

The electrons in the molecule are contained in molecular orbitals, which in the simplest form of the model are constructed from the atomic orbitals of the constituent atoms

The model correctly predicts relative bond strength, magnetism, and bond polarity

It correctly portrays electrons as being delocalized in polyatomic molecules ❯ The main disadvantage of the model is that it is difficult to apply qualitatively to polyatomic molecules

Molecular orbitals are classified in two ways: energy and shape

Energy

A bonding MO is lower in energy than the atomic orbitals from which it is constructed. Electrons in this type of MO are lower in energy in the molecule than in the separated atoms and thus favor molecule formation.

An antibonding MO is higher in energy than the atomic orbitals from which it is constructed. Electrons in this type of MO are higher in energy in the molecule than in the separated atoms and thus do not favor molecule formation.

Shape (symmetry)

Sigma σ MOs have their electron probability centered on a line passing through the nuclei

Pi π MOs have their electron probability concentrated above and below the line connecting the nuclei

Bond order is an index of bond strength

Bond order= # of bonding electrons - # of antibonding electrons/2

Molecules that require the concept of resonance in the localized electron model can be more accurately described by combining the localized electron and molecular orbital models.

The sigma σ bonds are localized

The pi π bonds are delocalized

Page 10

Chapter 17

Spontaneity, Entropy, and Free Energy

Page 11

Chapter 10

Liquids and solids

Condensed states of matter: liquids and solids

Held together by forces among the component molecules, atoms, or ions

Liquids exhibit properties such as surface tension, capillary action, and viscosity that depend on the forces among the components

Dipole-dipole forces

Attractions among molecules with dipole moments

Hydrogen bonding is a particularly strong form of dipole–dipole attraction

Occurs in molecules containing hydrogen bonded to a highly electronegative element such as nitrogen, oxygen, or fluorine

Produces unusually high boiling points

London dispersion forces

Caused by instantaneous dipoles that form in atoms or nonpolar molecules

Crystalline solids

Have a regular arrangement of components often represented as a lattice; the smallest repeating unit of the lattice is called the unit cell

Classified by the types of components:

Atomic solids (atoms)

Ionic solids (ions)

Molecular solids (molecules)

Arrangement of the components can be determined by X-ray analysis

Metals

Structure is modeled by assuming atoms to be uniform spheres

Closest packing

Hexagonal

Cubic

Metallic bonding can be described in terms of two models

Electron sea model: valence electrons circulate freely among the metal cations

Band model: electrons are assumed to occupy molecular orbitals

Conduction bands: closely spaced molecular orbitals with empty electron spaces

Alloys: mixtures with metallic properties

Substitutional

Interstitial

Network solids

Contain giant networks of atoms covalently bound together

Examples are diamond and graphite

Silicates are network solids containing SiO-O-OSi bridges that form the basis for many rocks, clays, and ceramics

Semiconductors

Very pure silicon is “doped” with other elements

n-type: doping atoms typically contain five valence electrons (one more than silicon)

p-type: doping elements typically contain three valence electrons

Modern electronics are based on devices with p–n junctions

Molecular solids

Components are discrete molecules

Intermolecular forces are typically weak, leading to relatively low boiling and melting points

Ionic solids

Components are ions

Interionic forces are relatively strong, leading to solids with high melting and boiling points

Many structures consist of closest packing of the larger ions with the smaller ions in tetrahedral or octahedral holes

Phase changes

The change from liquid to gas (vapor) is called vaporization or evaporation

Condensation is the reverse of vaporization

Equilibrium vapor pressure: the pressure that occurs over a liquid or solid in a closed system when the rate of evaporation equals the rate of condensation

Liquids whose components have high intermolecular forces have relatively low vapor pressures

Normal boiling point: the temperature at which the vapor pressure of a liquid equals one atmosphere

Normal melting point: the temperature at which a solid and its liquid have the same vapor pressure (at 1 atm external pressure)

Phase diagram

Shows what state exists at a given temperature and pressure in a closed system

Triple point: temperature at which all three phases exist simultaneously

Critical point: defined by the critical temperature and pressure

Critical temperature: the temperature above which the vapor cannot be liquefied no matter the applied pressure

Critical pressure: the pressure required to produce liquefaction at the critical temperature