Chapters 13 & 14

Energy: The capacity to do work

Potential Energy: stored energy due to position or condition. Chemicals can store energy; thus they have potential energy.

Kinetic Energy: energy in motion.

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Kinetic TheoryParticles have no attractive or repulsive forces

existing between the particles

Particles in gas move rapidly in constant motion. They travel in a straight path.

Total kinetic energy is conserved when particles collide.

Models of AtomsAtomic Models:Chemical properties of atoms, ions, and

molecules are related to the arrangement of the electrons within them.

John Dalton: 1st atomic model & considered the atom as a solid indivisible mass.

Dalton’s Atomic Theory1. All elements are composed of tiny indivisible

particles called atoms.

2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element.

3. Atoms of different elements can physically mix together or can chemically combine.

4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of element are never changed into atoms of another element as a result of a chemical reaction.

J.J. Thomson: revised Dalton’s model by proposing that electrons were stuck to the outside of the atom.

Ernest Rutherford: proposed the nuclear atom, in which electrons surround a dense nucleus composed of protons and neutrons.

Chadwick: discovered the neutron.

Discovery of the nucleusThis theory was discovered by Rutherford

who bombarded a sheet of gold foil with a beam of alpha particles surrounded by a fluorescent screen. They found that most of the particles passed through the foil, while a few were deflected.

Niels Bohr: student of Rutherford; proposed that electrons are arranged in concentric circular paths (orbits) around the nucleus.

Erwin Schrodinger: developed the quantum mechanical model

Atoms – tiny particles that make up matter

Structure of Atoms:Nucleus – center of the atom

protons – positively charged subatomic particles that is found in the nucleus; dictates the identity of the atom

-(Discovered by E. Goldstein using canal rays. Canal rays traveled from the positive metal plate to the negative metal plates)

Neutron: subatomic particle with no charge; found in the nucleus

- Discovered by James Chadwick

Electron: negatively charged; found outside the nucleus ( electron cloud)

- Discovered by J.J. Thomson using a cathode ray - the rays were attracted to a metal plate of positive charge.

ParticlesParticles SymbolSymbol ChargeCharge Relative Relative massmass

Actual Actual mass (g)mass (g)

ElectronElectron e- e- 1-1- 1/18401/1840 9.11x109.11x10-28-28

ProtonProton p+ p+ 1+1+ 11 1.67x101.67x10-24-24

NeutronNeutron nn00 00 11 1.67x101.67x10-24-24

Atomic Number - the # of protons in the nucleus

- the # of protons = the # of electrons

Mass Number – total # of protons & neutrons in an atom- To find the # of neutrons subtract the mass # from the atomic #

Example: Nitrogen (14

7N)

Mass number = 14Atomic number = 7

# of protons = 7# of electrons = 7# of neutrons = 14 – 7=

7

Isotopes - Has the same # of protons, but different #’s of neutrons

Atomic Mass Unit (AMU) – 1/12 the mass of carbon

Average atomic mass: weighted average of the masses of the isotopes of an element

In nature most elements occur as a mixture of two or more isotopes

Each isotope has a fixed mass and a natural percentage of abundance.

Average Atomic MassAvg atomic mass: what is found on the periodic table

=(mass)(% abundance) + (mass)(% abundance) +…

Practice ProblemAssume that element Uus is synthesized and

that it has the following stable isotopes:

284Uus (283.4 amu) 34.60% 285Uus (284.7 amu) 21.20% 288Uus (287.8 amu) 44.20%

What would the average atomic mass be?

Bohr’s ModelOrbits are known as energy levels.Electrons can move between energy levels.A quantum of energy is the amount of energy

required to move an electron up an energy level.

The higher the energy level the easier the electron can escape.

Bohr Models

Examples on the board

Elements #1-20

Atomic Orbitals1. Principal Quantum Numbers ( ) = 1,2,3,4….

2. Each principal level contains sublevels* Table 13.1 p. 364

3. Atomic orbitals are regions where electrons can be found. (Letter denotes the orbital)

S orbitals are spherical.

P orbitals are dumbbell-shaped.( exist in three different planes)

D orbitals have clover leaf shapes

F orbitals have complex shapes

4. The number & kinds of atomic orbitals depend on the energy sub level.

a. N = 1; 1 sublevel; 1s orbitalb. N = 2; 2 sublevels; 2s (1 orbital),

2p ( 3 orbitals)c. N = 3; 3 sublevels; 3s ( 1 orbital),

3p (3 orbitals), 3d (5 orbitals)d. N = 4; 4 sublevels; 4s (1 orbital),

4p (3 orbitals), 4d ( 5 orbitals), 4f (7 orbitals)

2n2

(n = principal quantum number).

This equals the maximum # of

electrons that the sublevel can hold.

Electron Arrangements in AtomsElectron Configurations:

1. Unstable systems tend to lose energy to become stable.2. Electrons try to form stable arrangements with the nucleus.3. The way in which electrons are arranged

around the nuclei of atoms is called electron configuration.

4. Three rules tell you how to find the electron configuration of atoms.a. Aufbau principle: electrons enter orbitals of lowest energy level first.

b. Pauli exclusion principle: and atomic orbital may describe at most two electrons. (arrows show the direction of electron spin)

c. Hund’s rule: when electrons occupy orbitals of equal energy, one electron enters each orbital, all of orbitals contain one electron with parallel spins.

Electron Configurations of Ions

5. When writing electron configurations for ions you must add or subtract the # of electrons gained or lost to create the ion.

Electron Configuration Practice

Elements #1-20

PERIODIC TABLE Periodic Table – an arrangement

of elements according to similarities in their properties

There are 92 naturally occurring elements.

Demitri Mendeleev – drew the first periodic table; Russian chemist arranged the first periodic table of elements in 1871. Arranged by atomic mass

* The periodic table contains chemical symbol, atomic number, & average atomic mass, physical state of each element, group numbers, and electron configuration.

Moseley: Later arranged the periodic table by atomic number. (Which is the one we use today.)

MODERN TABLEPeriods – horizontal rows (7 total)

Groups – vertical columns (has similar physical & chemical properties)

Metals – high electrical conductivity, luster, ductile, & malleable (Group 1 & 2A)

- Alkali Metals – Group 1A- Alkaline Earth Metals – Group 2A

Transition Metals & Inner Transition Metals – make up Group B (1B – 8B)

Nonmetals – poor conductors, non lustrous- Halogens – 7A- Noble Gases – 0

Metalloids – elements that border the stair step line

Group # = the outermost electrons

Periodic TrendsThe elements on the periodic table are arranged periodically so that trends can be recognized…

Trend of Ions1. You can determine the charge of an ion by what group it is in.1A = +1 5A = -32A = +2 6A = -23A = +3 7A = -14A = +/- 4

Trend of Electronegativity This refers to the ability of an atom to attract

the electrons of another atom to it.

Increases across the period ( left – right)

Decreases down the group ( top – bottom)

Trend of Electron affinityMeasure of the tendency for atoms to gain

electrons.

Increases across the period; this is caused by the filling of the valence shells

Decreases down the group; this is due to the electron entering an orbital far away from the nucleus

Trend of Ionization EnergyThe exact quantity of energy that it takes to

remove the outermost electron from the atom.

Factors affecting Ionization Energy: - nuclear charge - distance from the nucleus

Ionization energy increases across the period ( left – right) due to increased nuclear charge

Ionization energy decreases down the group ( top – bottom)

Trend of Atomic RadiusAtomic size is determined by how much

space the electron takes up. It is also depends on how far its valence electrons are from the nucleus.

The atom will be large if the electron is far from the nucleus

- size increases down a group (top – bottom)

The atom will be small if the electron is close to the nucleus

- size decreases across the period ( left – right)

This is due to an increase in nuclear charge pulling them closer… the energy level stays the same

Trend of Metallic/Non-Metallic PropertiesMetallic properties: elements will form

cations as they lose electrons (+ve charge)

Non-Metallic properties: elements form anions as they gain electrons (-ve charge)

Trend of Melting / Boiling PointsMelting and Boiling point increase from the

right side of the periodic table until it reaches aluminum and silicon

Here, melting point and boiling point then begin to decrease.

Trend of ReactivityHow likely/vigorously an atom is to react with

other substancesMetals:- Period: decreases from left to right- Group: increases down the group

The farther left and down you go the easier it is for electrons to be taken away. (Higher Reactivity)

Trend of ReactivityNon-Metals- Period: increases from left to right- Group: decreases down the group

The farther right and up you go the higher electronegativity – vigorous exchange of electrons

Classification of ElementsElements can be classified into 4 groups based

on electrons.1. Noble gases: outermost s & p sublevels are filled. Belong to group 0. (Also called

inert gases.)2. Representative elements: outermost s or p sublevel is partially filled

3. Transition metals: metallic elements in which the outermost s sublevel and near d sublevel contain electrons. (Group B elements)

4. Inner transition metals: metallic elements in which the outermost s sublevel and nearby f sublevel generally contain electrons. (Lanthanide & Actinide series)

Light and Atomic SpectraLight consists of electromagnetic waves.Light has a velocity of 3.0 x 10 8 m/s.Amplitude: is the wave height from origin to

crest.Wavelength (λ): distance between crest.Frequency (ν): number of wavelength to

pass a given point per unit of time. (units = hertz Hz)

c = speed of light (3.00 x 10 8 m/s)λ= wavelengthν= frequency

c=λν

Example: Calculate the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is 5.10 x 10 14 Hz (5.10 x 10 14 s-1).

c = 3.00 x 108 m/sFrequency (ν) = 5.10 x 1014 s-1

wavelength (λ) = ??? m

Frequency & wavelength are inversely related.

Electromagnetic spectrum: series of waves at different wavelengths (radio waves, radar, microwaves, infrared, visible light, ultraviolet, x-rays, gamma rays, cosmic rays)

Every element emits light when it is excited by the passage of an electric discharge through its gas or vapor.

Black & White LightBlack light – All colors absorbed

White Light – All colors reflected

What happens for you to see colors?

Planck’s constant (h)– 6.63 x 10 -34 J x sE = energyh = Planck’s constantν = frequency

E = h x ν

Example: Calculate the energy (J) of a quantum of radiant energy (the energy of a photon) with a frequency of 5.00 x 1015 s-1.

ν = 5.00 x 1015 s-1

h = 6.63 x 10 -34 J x sEnergy(E) = ??? J