AP ChemistryWinter Break Assignment

Name Date

This assignment will encompass the concepts concerning electronic structure, Bohr model of the atom, quantum numbers, orbital shapes, sizes, and hybridization, the Quantum Mechanical model of the atom, electron configuration, periodic trends, lewis structures, molecular geometry and polarity, phase diagrams, intermolecular forces, and crystal structure of molecular and ionic solids. You should know and understand the concepts behind these topics upon the return to school on January 2, 2013. Calculations will be reviewed before a graded assessment is assigned, but the basic concepts will not be addressed in depth and it is your responsibility to study the material provided, along with reading Chapters 6, 7, and 9 of the text. Problem sets listed below will be due on Tuesday, January 8, 2013. Complete all worksheets provided and attach the problem set to this packet.

Chapter 6: page 158- Summary Problem, #1, 2, 3, 6, 7, 8, 10, 11, 14, 19, 27, 29, 31, 39, 50, 61, 67, 71, 73, 76, 80

Chapter 7: page 191-Summary Problem, #17, 19, 21, 24, 27, 29, 37, 47, 55, 57, 63, 69, 81

Chapter 9: page 252- Summary Problem, #1, 3, 5, 7, 9, 11, 17, 21, 31, 33, 45, 51, 56, 60, 61, 66, 74, 76

Chapter 6: Electronic Structure and the Periodic Table

6.1: Light, Photon Energies, and Atomic Spectra Structure of the Atom: the structure of the atom has changed drastically over time,

beginning with Democritus’ idea of the atom being an indivisible sphere and continuing to evolve today.

o 450 B.C. Democritus

All matter is made up of “atomos” (indivisible particles) At a certain point, matter could not be broken down any

further Aristotle

Disagreed with Democritus 4 Elements: Earth, Wind, Fire, Water

Both were philosophers, not scientists. Aristotle’s ideas were believed for nearly 2000 years.

o Late 1700s: 3 important discoveries Antoine Lavoisier: Law of Conservation of Matter Joseph Proust: Law of Definite Proportions Ben Franklin: Electricity

Objects could have 1 of 2 electrical charges (+/-) Like charges repel, opposites attract Some objects readily pick up electrical charge (static

electricity, metals)o 1803

John Dalton and his Atomic Theory Each element is composed of small particles (atoms) All atoms of the same element are identical, but differ from

those of any other element Atoms are neither created or destroyed in a chemical reaction A given compound always has the same relative number and

kinds of atoms Atoms are small, solid, spherical, indivisible, neutral MARBLE MODEL

o 1840 Michael Faraday

Atomic Structure is related to Electricity! Used Cathode Ray Tubes and discovered a stream of particles **Cathode Ray tube is a partially evacuated glass tube

containing a gas at a low pressure**o 1869

Dmitri Mendeleev and Lothar Meyer published almost identical periodic tables. (Mendeleev got more credit than Meyer)

o 1871 Mendeleev predicted unknown elements and their properties.

o 1896 JJ Thomson

Used Cathode Ray Tube to study electric discharge in a vacuum Used ANODE with a hole to allow a ray to pass through Magnets and Electricity can alter the ray’s path in a

mathematically predictable way Concluded the cathode ray was NEGATIVE PARTICLES coming

from the cathode (ELECTRONS) **Proved Democritus wrong (Atoms are divisible) Charge to mass ratio = 1.76 x 108 C/g PLUM PUDDING MODEL/CHOCOLATE CHIP COOKIE MODEL

Henri Becquerel Radioactivity: Spontaneous emission of radiation Marie Curie and Pierre Curie followed in his footsteps

o 1900 MAX PLANCK

Proposed that there was a predictable relationship between a QUANTUM of energy and the frequency of radiation

An object emits energy in small, specific amounts (QUANTA) E = h h = Planck’s Constant = 6.626 x 10-34Js Planck’s suggestions were thought to be extremely radical at

the time 1905

o ALBERT EINSTEIN Expanded on Planck’s theory by stating that electromagnetic radiation

(EMR) has a dual wave-particle nature Each individual particle of light can carry 1 quantum of energy PHOTON: a particle of EMR having zero mass and carrying a quantum

of energy Ephoton = h Minimum energy ~Minimum frequency Must reach this minimum energy before a photon will be released

o 1909 Robert Millikan

Oil Drop Experiment: Spray drops in apparatus, give drops negative charge using x-rays

Varying the charge effected droplets’ rate of fall Determined charge of electron: 1.60 x 10-19C Mass of electron = 9.11 x 10-28g

o 1911 Ernest Rutherford

Gold Foil Experiment: Bombarded a thin sheet of gold with alpha particles and most of them passed through. A few bounced back

Discovery of NUCLEUS

Suggested that electrons orbit nucleus, but could not explain his findings.

Nucleus: Densely packed bundle of matter with a positive charge

Henry Moseley Elements fit into better patterns when arranged by increasing

ATOMIC NUMBER (# of protons) PERIODIC LAW : The physical and chemical properties of the

elements are periodic functions of their atomic numbers. PERIODIC TABLE : An arrangement of the elements in order of

atomic number so that elements with similar properties (PERIODICITY) are in the same GROUP.

PERIODICITY: properties repeat themselves at regular intervals throughout the periodic table.

o 1913 Niels Bohr

He developed a model of the atom that linked ELECTRONS with PHOTON EMISSION

Electrons can circle the nucleus in specific paths (ORBITS) and when in one of these orbits, the atom has a DEFINITE and FIXED ENERGY. Within each orbit, their exists SUBLEVELS.

o Orbit closest to the nucleus = lower energy stateo Orbits further from nucleus = higher energy stateo Electrons CANNOT EXIST between orbits (energy

levels) PLANETARY/SOLAR SYSTEM MODEL

*Bohr’s idea explained only Hydrogen, not atoms with more than 1 electron*

o QUANTUM MECHANICAL MODEL: Differs from the Bohr Model in that The Kinetic Energy of an electron is inversely related to the volume of

the region to which it is confined

It is impossible to specify the precise position of an electron in an atom at a given instant

o 1924: LOUIS DE BROGLIE Bohr’s quantized electrons behaved like waves Duality of matter: Electrons could be considered as waves confined to

a certain space around an atomic nucleus. (certain space = specific frequency = specific energy)

Experimentation proved that:1. Electrons (like light waves) could be bent of DIFFRACTED

a. Diffraction: The bending of a wave as it passes by the edge of an object (ex: edge of an atom in a crystal)

2. Electron beams (like waves) can INTERFERE with each othera. Interference: Waves overlap which results in a reduction of energy

in some areas and an increase of energy in other areas

o HEISENBERG UNCERTAINTY PRINCIPLE Detection of electrons by observing their interaction with photons Photons have about the same energy as electrons and when used to

detect electrons the photon will change the electron’s path. This results in UNCERTAINTY when trying to locate electrons

HEISENBERG UNCERTAINTY PRINCIPLE states: it is impossible to determine simultaneously both the POSITION and VELOCITY of an electron (or any particle)

We must rely on PROBABILITY that an electron is in a certain location and moving at a certain speed.

Photons are shot at an electron and when they collide we are able to find the exact LOCATION of an electron, but this collision causes the VELOCITY to change.

o SCHRODINGER WAVE EQUATION Schrodinger used the dual wave-particle hypothesis to develop an

equation that treated electrons in atoms as waves Naturally accounted for Bohr’s idea of quantization Together with HUP, laid a foundation for modern QUANTUM THEORY:

Describes mathematically the wave properties of electrons and other small particles

Based on HUP, we can only find the probability of locating an electron at a given point around the nucleus

Electrons did not exist in neat paths (BOHR’S ORBITS) but instead in certain regions (ORBITALS: 3-D region around the nucleus that indicates the probable location of an electron)

o PAULI EXCLUSION PRINCIPLE: No 2 electrons can be in the same exact place at the same exact time.

ELECTROMAGNETIC SPECTRUM Made up of all forms of ELECTROMAGNETIC RADIATION All move at a constant speed of 2.998 x 108 m/s (167,000,000 mi/hr)

WAVELENGTH: ( : lambda) The distance between corresponding points on adjacent waves

o 1 nm = 1 x 10-9mo This value must always be converted to METERS

FREQUENCY: ( : nu) The number of waves that pass a given point in a specific time

o 1 wave/second = 1Hertz (Hz) c = Amplitude: ( Ψ: psi) Height of crest/trough

Photon Energieso Light is a stream of tiny particles (photons) and the energy can be calculated

using the following equationo E = h

Energy is measured in Joules h = Planck’s Constant = 6.626 x 10-34Js

Atomic Spectra: All species emit photons at specific, predictable wavelengths. These wavelengths (and their corresponding frequencies and energies) can be studied to determine the identity of the species.

o The simplest of the atomic spectra is HYDROGEN. Hydrogen’s atomic spectra has been studied extensively in an effort to understand more about electronic behavior. Hydrogen emits groups of photons in 3 different regions of the electromagnetic spectrum:

Lyman Series (UV region) Balmer Series (Visible region) Paschen Series (Infrared region)

6.3: Quantum Numbers1. PRINCIPAL QUANTUM NUMBER (n)

a. Main energy level of electronb. Positive whole numbers (1,2,3,…)c. Energy increases as distance from the nucleus increasesd. More than 1 electron can have the same n-value

e. Total number of orbitals in a given shell = n2

2. ANGULAR MOMENTUM QUANTUM NUMBER (l)a. Indicates the sublevel, which determines the general shape

of the electron cloudb. This value ranges from 0 – (n-1)c. l = 0 = s-sublevel

1 = p-sublevel 2 = d-sublevel 3 = f-sublevel

Shapes of orbitals:

ENERGY LEVEL

SUBLEVEL # of ORBITALS # of ELECTRONS

1 s 1 2

2 p 3 6

3 d 5 10

4 f 7 14

3. MAGNETIC QUANTUM NUMBER (ml)a. Designates the direction in space in which the orbital is

orientedb. More than 1 electron can have the same SHAPE, but a

different ORIENTATIONc. –l to +l

4. SPIN QUANTUM NUMBER (ms)a. TWO possible values (+/- ½)b. TWO fundamental spin states of an electron in an orbitalc. A single orbital can hold a maximum of 2 electrons with

opposing spins

6.5: Electron Configurations in Atoms

ELECTRON CONFIGURATIONSo The arrangement of electrons in an atom o Distinct address for each elemento Electrons will assume arrangements with lowest

o possible energyo RULES:

1. ORDER of electrons occupying each orbitala. Aufbau Principle: an electron occupies the lowest energy

level/orbital that can receive it2. Importance of SPIN

a. Pauli Exclusion Principle: no 2 electrons in the same atom can have the same set of 4 QNs

3. Electrons will occupy ALL orbitals before pairing in an orbitala. Hund’s Rule: orbitals of equal energy are each occupied by

a second electron and all electrons in singly occupied orbitals must have the same spin

o 3 METHODS OF NOTATION1. Orbital Notation

a. Unoccupied orbital represented by a line: b. Orbital containing one electron: c. Orbital containing 2 electrons: (opposing spins)d. Each line (orbital) is labeled with Principal Quantum

Number2. Electron-Configuration Notation

a. Eliminates lines and arrows of orbital notationb. Electrons in each sublevel indicated by adding a

SUPERSCRIPT to the sublevel3. Noble-Gas Notation

a. Beginning with 3rd period, electron-configuration notation can be abbreviated using the noble gases

6.8: Periodic Trends

1. ATOMIC RADIUS ½ the distance between the nuclei of 2 identical atoms that are

bonded together DECREASES from left to right INCREASES from top to bottom

2. IONIC RADIUS

Formation of a CATION always leads to a DECREASE in atomic radius Formation of an ANION always leads to an INCREASE in atomic radius

3. IONIZATION ENERGY An ion is an element/group of elements that has a positive or negative

charge Ionization is any process that results in the formation of an ion The energy required to remove one electron from a neutral atom of an

element. Those elements that EASILY lose electrons are HIGHLY REACTIVE INCREASES from left to right (due to increasing nuclear charge) DECREASES from top to bottom (due to increasing size of electron

cloud) Electrons further from the nucleus are more easily removed SUCCESSIVE IONIZATION ENERGY: Removal of electrons from

cations. Each successive electron removed from an ion feels an increasingly stronger effective nuclear charge.

Noble Gases and Ions with Noble Gas Electron Configuration have extreme stability making it very difficult to lose more electrons

4. ELECTRONEGATIVITY A measure of the ability of an atom in a chemical compound to

ATTRACT electrons INCREASES from left to right DECREASES from top to bottom

5. ELECTRON AFFINITY

Energy change that occurs when a neutral atom acquires an electron INCREASING or DECREASING (no real pattern) DECREASES from top to bottom

Chapter 7: Covalent Bonding

7.1: Lewis Structures and the Octet Rule

1916: Gilbert Newton Lewis suggested that noble gases are very unreactive and that nonmetal ions can achieve this stability by sharing electrons.

The Octet Rule states that the most stable arrangement of electrons consists of 8 valence electrons, resembling the noble gases.

o Exceptions to the Octet Rule Hydrogen and Helium Boron: generally forms 3 bonds because it has only 3 valence

electrons Elements bonded with F/O/Cl (Highly Electronegative) sometimes

hold more than 8 valence electrons; an expanded octet or expanded valence shell

Lewis Structureso Atomic Symbols represent nuclei and inner-shell electrons.o Dots/Dashes between atoms represent SHARED PAIR of electrons.o Dots next to only one atom represent LONE PAIR of electrons.o Structural Formula = only shared pairs between atoms are represented,

usually by dashes but can also be represented by lines.o RULES

1. Count the # of valence electrons2. Write symbols of LEAST ELECTRONEGATIVE elements in the MIDDLE3. Assign dots as valence electrons around all elements. Make sure not to

exceed the total number of valence electrons.4. If each element in the structure does NOT have 8 valence electrons,

rearrange to form double or triple bonds. If there are electrons left over, add electron pairs to central atom. (This usually only happens when the central atom is bonded to HIGHLY ELECTRONEGATIVE elements)

o Multiple Covalent Bonds Double bond: 2 shared pairs between 2 atoms. Triple bond: 3 shared pairs between 2 atoms.

o Resonance Structures: Bonding cannot be accurately represented by a single Lewis Structure

Ex: O3 (ozone)O = 6(3) = 18ve

o Ions: atom/group of atoms that has a charge (+) ion (cation): subtract charge from total number of valence

electrons (-) ion (anion): add charge from total number of valence electrons

7.2: Molecular Geometry

VSEPR Theory: Valence-shell electron-pair repulsion theoryo Electrons repel each other, causing bonds to be as far away from one another

as possible.o Most important atom when determining geometry is the CENTRAL ATOMo ABE Notation: describes how electrons are shared around the central atom.

A = Central Atom B = # of atoms bonded to central atom E = # of lone pairs on central atom

SHAPE # ATOMS BONDED TO CENTRAL (B)

# LONE PAIRS ON CENTRAL (E)

ABE NOTATION

Linear 12

3,1,00

ABE3/ABE/ABE2/ABAB2

Bent/Angular 22

12

AB2EAB2E2

Trigonal Planar 3 0 AB3

Tetrahedral 4 0 AB4

Trigonal Pyramidal 3 1 AB3ETrigonal

Bipyramidal5 0 AB5

Octahedral 6 0 AB6

7.3: Polarity of Molecules Intermolecular Forces (IMF): Forces of attraction between MOLECULES

o Weaker than bonds (covalent/ionic/metallic) holding atoms together.o As temperature increases, KE increases, and molecules begin to pull away

from one another.o IMFs are between MOLECULES (Covalently bonded groups of atoms)o 3 Types:

1. London Dispersion Forces Results from the constant motion of electrons and creates

instantaneous dipoles Present in ALL ATOMS and MOLECULES Noble Gases and Nonpolar Molecules only have this type of IMF Increase with increasing # of electrons and increasing mass

2. Dipole-Dipole Forces Dipole: created by equal but opposite charges separated by a

distance (bond length) Ex: H2O

= Delta + = dipole; (+) near less electronegative element, near more

electronegative element Forces of attraction between POLAR MOLECULES POLARITY depends on both ELECTRONEGATIVITY difference as

well as GEOMETRYGEOMETRY ALWAYS NONPOLAR ALWAYS POLAR SOMETIMES POLAR

Linear Bent/Angular

Trigonal Planar Tetrahedral

Trigonal Pyramidal Trigonal

Bipyramidal

Octrahedral

Dipole-Induced Dipole: A polar molecule can INDUCE a dipole in a nearby NONPOLAR molecule by temporarily attracting its electrons (This force is WEAKER than a Dipole-Dipole force)

3. Hydrogen Bonding Hydrogen is bonded to the following highly electronegative

elementsi. NITROGEN

ii. OXYGENiii. FLOURINE

Strongest IMF Ex: H2O

7.4: Atomic Orbitals and Hybridization

sp hybrid orbitals: Bonding electrons located in the s-orbital of 1 atom and the p-orbital of a second atom will for 2 sp-hybrid orbitals

o s + p 2sp

sp2 hybrid orbitals: s + 2p 3sp2

sp3 hybrid orbitals: 1s + 3p 4sp3

sp3d hybrid orbitals: 1s + 3p + 1d 5sp3d

sp3d2 hybrid orbitals: 1s + 3p + 2d 6sp3d2

Extra electrons in a multiple bond are NOT located in hybrid orbitals. Only 1 of the shared pairs in a multiple bond exists in the hybrid orbital.

o Sigma Bonds ( ): Consist of an electron pair occupying a hybrid orbital σ

o Pi Bonds ( ): Consist of an electron pair occupying the unhybridized orbitalπ SINGLE BOND: 1 sigma DOUBLE BOND: 1 sigma and 1 pi TRIPLE BOND: 1 sigma and 2 pi

Chapter 9: Liquids and Solids

9.2: Phase Diagrams

9.4: Network Covalent, Ionic, and Metallic Solids

1. Network Covalent Solids: a continuous network of covalent bonds joins atoms.

High Melting Point Insoluble in all common solvents Poor electrical conductors

2. Ionic Solids: Consist of cations and anions

Nonvolatile with high melting points Do not conduct electricity unless dissolved or in the liquid state because as solids,

ions are in fixed, permanent positions Many are soluble in water Strength of ionic bonds can be estimated using Coulomb’s Law and depends on 2

factorso Charges of the ionso Size of the ions

3. Metals: Consist of a sea of electrons

High electrical conductivity High thermal conductivity Ductility and Malleability Luster Insolubility in water and other common solvents Mercury can dissolve other metals to form a solution called an amalgam.

9.5: Crystal Structures**Table 9-5: Structures and Properties of Types of Substances (p.245)**A unit cell is used to represent the structural geometry of a solid

Metals: 3 types of solid structureso Simple Cubic Cell (SC): 8 atomso Face-Centered Cubic Cell (FCC): 9 atomso Body-Centered Cubic Cell (BCC): 12 atoms

**Table 9-6: Properties of Cubic Unit Cells**

Ionic Crystals: 3 types of lattices (see page 249 for diagrams of ionic crystals)o Face-Centered Cubic Lattice-1 (FCC1)o Face-Centered Cubic Lattice-2 (FCC2)o Body-Centered Cubic Lattice (BCC)

Wavelength, Frequency, and Energy Calculations

1. Green light has a wavelength of 5.0 x 102 nm. a. What is the energy, in joules, of one photon of green light?b. What is the energy, in joules, of 1.0 mol of photons of green light?

2. Violet light has a wavelength of about 410 nm. a. What is the frequency?b. What is the energy of one photon of violet light?c. What is the energy of 1.0 mol of violet light?

3. The most prominent line in the spectrum of aluminum is 396.15 nm. a. What is the frequency of this line?b. What is the energy of one photon with this wavelength?

4. The most prominent line in the spectrum of magnesium is 285.2 nm. Other lines are found at 383.8 nm and 518.4 nm.

a. In what region of the electromagnetic spectrum are these lines found?b. Which is the most energetic line?c. What is the energy of 1.00 mol of photons with the wavelength of the MOST

ENERGETIC line?

5. The light from an amber signal has a wavelength of 595 nm, and that from a green signal has a wavelength of 500.0 nm.

a. Which has a higher frequency?b. Calculate the frequency of amber light.

6. You are an engineer designing a switch that works by the photoelectric effect. The metal you wish to use in your device requires 6.70 x 10-19J/atom to remove the energy from a single electron. Will the switch work if the light falling on the metal has a wavelength of 540nm or greater? Why or why not?

7. If a large pickle is attached to two electrodes with a 110-Volt power supply, the pickle begins to glow with a yellow color. Knowing that pickles are made by soaking cucumbers in a concentrated salt (NaCl) solution describe why the pickle might emit yellow light when electrical energy is added.

Quantum Numbers and Electron Configuration

1. Determine the 4 quantum numbers for the following elementsa. Nb. Mgc. Ard. Fee. Rbf. Br

2. What are 2 exceptions to electron configurations and the Aufbau Principle, and why do these exceptions occur?

3. What are the 4 Quantum Numbers for the following electrons?a. 4d3

b. 6s1

c. 2p5

4. For each of the following elements, write the Orbital Notation, Electron Configuration Notation, and Noble Gas Notation on a separate sheet of paper.

a. Li b. Mgc. Gd. Ce. Pf. Seg. Fh. Kri. Crj. Ag

5. When n = 4, what are the possible values of l?6. When l = 2, what are the possible values for ml?7. In a 4s orbital, what are the possible values of n, l, ml and ms?8. When n = 4, l = 2, and ml = -1, what orbital does this refer to? (Ex: 1s)9. Why is the following set of quantum numbers not possible for an electron in an

atom?n = 2, l = 2, ml = 0

10. Which of the following orbitals cannot exist according to the quantum theory: 2s, 2d, 3p, 3f, 4f, and 5s. Explain your answer.

11. A given orbital has a magnetic quantum number of ml = -1. Which sublevel is not possible? (Ex: s, p, d, or f)

Periodic Trends

1. Which has a larger atomic radius: Sodium or Cesium? Explain your reasoning.

2. Which has a smaller radius: Calcium atom or Calcium +2 ion? Explain your reasoning.

3. Which has a higher 1st ionization energy: Lithium or Oxygen? Explain your reasoning.

4. Which ionization energy for Radium is the highest and explain your answer: a. 1st

b. 2nd

c. 3rd

d. 4th

5. Which has a lower electronegativity value: Silicon or Sulfur? Explain your reasoning.

1. Which has a larger atomic radius: Aluminum or Silicon

2. Which has a smaller atomic radius: Strontium or Rubidium

3. Which has a larger ionic radius: Fluorine atom or Fluorine ion

4. Which has a smaller ionic radius: Calcium atom or Calcium ion

5. Which has a larger 1st ionization energy: Lithium or Neon

6. Which has a larger 2nd ionization energy: Phosphorus or Sulfur

7. Which has a higher electronegativity: Chlorine or Iodine

8. Which has a lower electrongativity: Magnesium or Barium

9. Which ionization energy for Gallium is the highest: 1st, 2nd, 3rd, or 4th

Lewis Structures, Molecular Geometry, Polarity, and Intermolecular Forces

DIRECTIONS: For each of the following compounds,a) Draw Lewis Structures (be sure to indicate the number of valence electrons)b) Identify the ABE notation c) Predict the geometry for each of the moleculesd) Predict the polarity of the molecule (If polar, show direction of polarity)e) State the strongest IMF present

1. AsCl5

2. SeF6

3. PH3

4. SF2

5. C2HI6. BCl3

7. CH3CHCCHCH3

8. HOCl9. CHF3

10. PCl5

11. SeF6

12. PH3

13. Cl2CO14. SiO2

15. CH3CH2OCH3

16. CH3CH2CH2CH2NH2

17. HI18. H2Se19. OClOH20. CH4

21. (CO3)2—

22. NH4+

23. C2H6

24. CO2

25. HCN26. C3H8

27. C3H6

28. BH3

29. SeH6

30. F2

31. CH3COOH