AP Chem

Unit 1 – Atomic Structure and Properties

1.1 - Moles and Molar Mass1.2 – Mass Spectroscopy of the Elements

1.3 – Elemental Composition of Pure Substances1.4 – Composition of Mixtures

1.5 – Atomic Structure and Electron Configuration

1.6 – Photoelectron Spectroscopy1.7 – Periodic Trends

1.8 – Valence Electrons and Ionic Compounds

Moles and Molar Mass

This is a review topic from SCH3U.

Exercise:Given the following scenarios:

A. 2.0 x 1023 molecules of glucose, C6H12O6

B.X g calcium carbonate, CaCO3 (sample given)

C.0.25 mol propane, C3H8

Rank in order of increasing:

a. Mass

b. Mol

c. Number of particles

d. Number of atoms

e. Volume

Example AP MC questions:

Mass Spectroscopy of the Elements

Here is a mass spectrum of the element zirconium. It shows the relative abundance of each isotope, and can be used to determine the average atomic mass of this element.

Average atomic mass =

Example: Determine the average atomic mass of each of the following and identify the element.

How does a mass spectrometer work?

A sample of the substance is injected into the machine. The sample is then ionized as only positively charged cations are read by the detector. A common type of ionization method is to use a high-energy electron beam which knocks off a valence electron to form a cation. The cations are then accelerated toward a mass filter which separates them by mass. This is done using a magnet. Particles with less mass are deflected ___________ by the magnet, while particles with more mass are deflected _____________. The separated particles are caught by a detector which detects and quantifies the separated ions.

In this course, we only look at mass spectra of atoms, but mass spectrometry is a valuable tool for many areas of research since more complicated molecules can also be analyzed with it.

Example AP MC questions:

Elemental Composition of Pure Substances

Molecules are help together by ______________________ bonds exclusively.

For example:

Formula units are used to describe ___________________ compounds. _____________ bonds are electrostatic forces of attraction between oppositely charged particles. ________________ ions, both cations and anions, have ____________ bonds holding that ion together.

For example:

Law of Definite Proportions:

A pure sample of a compound will always have the same ratio of masses of each element. For example, the mass of a pure sample of MgCO3 is 29% magnesium, 14% carbon and 47% oxygen. If one was given sample of MgCO3 that did not match up with these values, you would know it was impure.

Example AP MC questions:

Empirical and Molecular Formulas:

Method to determine an empirical formula (and molecular formula) given % composition:

1. Change % to grams (if necessary)

2. Use molar mass of element to determine number of mol

3. Place mol as subscript after each element and simplify to whole number by dividing all by lowest value. Remember that one can only round if decimal is less than 0.2 or greater than 0.8. [hint: if near 0.2 x5, 0.25 x4, 0.33 x3, 0.5 x2, 0.67 x3, 0.75 x3, 0.8 x5)

4. To determine a molecular formula, divide the molar (or molecular mass) by the empirical mass to determine the number of empirical units in the molecule.

Example: Determine the molecular formula of fructose if its molar mass is 180.2 g/mol and has a composition 40.00% C, 6.72% H, and 53.29% O.

Example AP MC question:

Composition of Mixtures

Many samples of chemicals are not pure. We can define percent purity as:

mass of pure compound in the impure sample x 100%

total mass of impure sample

Looking at percentage composition, it is possible to assess the purity of a sample based at the theoretical composition of the pure substance.

For this topic, you will also be completing the Gravimetric Analysis (formative) lab (after the AP MC questions)

Example AP MC questions:

AP Chem

Gravimetric Analysis – Finding the amount of chloride in an unknown sample

In this lab, you will be using a technique called gravimetric analysis. Gravimetric analysis is an example of quantitative analysis. The amount of substance in a sample is determined through finding out how much of a known _______________ it can form during a reaction. In this experiment, we will find the amount (mass percent) of chloride ion in an unknown chloride salt and from there determine its identity. To do this you will:

1. Determine an appropriate compound to react with the unknown salt. Ms. Sardone will prepare a solution of it to use in your experiment.

2. Carry out an experiment that will let you determine how much chloride was present in the sample.

3. Use calculations of mass percent to determine the identity of the unknown salt.

Possible Identities of the Unknown:AlCl3, NaCl, PbCl2, KCl

Part 1 – What reagent will you use and why?

Things to consider:

· What type of compound is your unknown?

· Is it soluble in water?

· What types of reactions does it undergo?

· What conditions are needed to carry out this type of reaction?

· What types of characteristics are needed in the product of my reaction so that it can be analyzed?

Procedure

1. Weigh out approximately 0.2-0.4g of the unknown chloride salt into a 250mL beaker. Record the exact amount you are using.

2. Dissolve the salt using approximately 150mL of distilled water. Add approximately 1mL of 6M HNO3 and stir.

3. Add 20 mL of 0.5M _______________ solution. Heat the solution gently for 10 minutes and then put beaker in an ice bath for another 10 minutes.

4. Find the mass of a filter paper and filter the product to remove the solvent. Wash with distilled water a few times and once with acetone.

5. Carefully remove the filter paper from the funnel and place on a watch glass. Leave in a cool, dark location to dry further (overnight).

6. Find and record the mass of the product and filter paper.

Mass of unknown chloride salt

Mass of filter paper

Mass of filter paper and silver chloride

Mass of silver chloride

Follow-up questions

1. Write the net ionic equation for this reaction.

2. Determine:

· the moles of silver chloride produced

· the moles and mass of chloride

· the percentage of by mass of the unknown that is chloride

3. Which of the following chloride salts is most likely the unknown?

AlCl3 or NaCl or PbCl2 or KCl

Atomic Structure and Electron Configuration

The Bohr-Rutherford model of the atom, while easy to understand and use, is not the current model of the atom. Developed shortly after the Bohr model, the quantum mechanical model has been the accepted model. However, due to its abstract nature, it is studied in later chemistry courses.

We will:

1. look at the historical development of this model

2. use it to show electron arrangements called electron configurations

3. model the electron configuration by constructing orbital diagrams

4. go into the specifics of the model to explain some chemical and physical phenomena of the elements as well as periodic trends

Theories of Atoms and their Structure

Scientist

Important Dates

Details of their atomic model

Diagram

Limitations

John Dalton

1808

· indestructible sphere

· atoms combine to form other substances

· did not explain why atoms

combine in the ratios in which they do

J. J. Thompson

1897

· used gas discharge tube

· electrons

· uniform positive charge over sphere

· “plum pudding” model

· gold foil experiment (later)

Ernest Rutherford

1911

· gold foil experiment

· nucleus with positive charge

· McGill

· Nobel Prize

· didn’t explain total mass of the atom

· didn’t explain why electrons don’t collapse into the nucleus

Neils Bohr

1913

· electrons in specific energy levels

· fixed circular orbits

· only explained line spectra of hydrogen, no other atoms

Henri DeBroglie

1920s

· if light has properties of matter, then matter or particles have properties of light

Erwin Schrodinger

1926

· wave equation

· quantum mechanical model

James Chadwick

1932

· neutrons in nucleus

How to write electron configurations and construct orbital diagrams

Orbitals: 3-D regions around the nucleus where the probability of finding an electron is greatest

Each orbital can hold ____ electrons

There are different types of orbitals with different shapes:

Type of orbital

Shape

Number of these orbitals at each energy level

Total number of electrons in that type of orbital at each energy level

s

p

d

f

Too complicated to draw!

Electron Configurations:

For example:S has the electron configuration:1s22s22p63s23p4or [Ne] 3s23p4

V has the electron configuration:1s22s22p63s23p64s23d3or [Ar] 4s23d3

Try these:

element

electron configuration

Ca

Se

F

Zr

Orbital Diagrams:

· A box is used to represent each orbital. S-orbitals are singular, p-orbitals in groups of 3, d-orbitals in groups of 5, f-orbitals in groups of 7. Always label the orbitals.

· Electrons are represented by arrows. Two electrons in the same orbital will have arrows in opposite directions. Only pair up electrons.

For example:F has the electron configuration and orbital diagram:

Try these:

C

P

V

Quantum Mechanical Model of the Atom - Theory

What is an orbital?

Quantum Numbers

Each atomic orbital is identified by three quantum numbers. They are:

1.

2.

3.

Each electron in an orbital is also given a ______ quantum number, ms.

Principle Quantum Number (n):

· n =

· is related to the _____________ and ________________ of the orbital

· as the number increases, the orbital becomes ____________ and the electrons spend more time further from the nucleus

· maximum number of orbitals at that energy level =

· maximum number of electrons at that energy level =

Angular Momentum (Orbital Shape) Quantum Number (l):

· Identifies the ________________ of the orbital

· l =

· when n = 1, l =

· when n = 2, l =

· when n = 3, l =

etc.

· The values of l are most commonly described by letters:

Value of l

0

1

2

3

4

5

Letter

· These letters indicate the _____________ of the orbitals:

s – orbitals:

· s-orbitals are _________________ shaped

· Each s-orbital can hold _____________ electrons

p – orbitals:

It is important to realize that orbitals are solutions to mathematical functions. They describe the probability of finding the electron in a particular spot. The spheres and dumbbell diagrams seem concrete and solid, so it can be easy to think of the orbitals as containers. Remember, orbitals do not have physical properties – they exist in the imagination.

· Start at the ___________ energy level

· There are ________ p-orbitals at each energy level

· p-orbitals are _________________ shaped

· Each p-orbital can hold _____________ electrons

d – orbitals:

· Start at the ___________ energy level

· There are _______ different shapes

· Each d-orbital can hold _______ electrons

f- orbitals:

· There are 7 different shapes

· Each f-orbital can hold 2 electrons

Magnetic Quantum Number (ml)

· Indicates the orbital’s spatial _____________________ around the nucleus

· mll =

· when:n = 1l = 0 ml = 0

n = 2l = 0 ml =

l = 1ml =

ml =

ml =

n = 3

Spin Quantum Number (ms)

· describes the electron _________ since electrons behave as tiny ______________ spinning on their axis in either ____________________ or ____________________ directon.

· ms =

Using quantum numbers:

· We can assign 4 quantum numbers to each of the electrons in all of the atoms on the periodic table

· For example:

· Lithium has 3 electrons:

n = 1 l = 0 ml = 0 ms = + ½

n = 1 l = 0 ml = 0 ms = - ½

n = 2 l = 0 ml = 0 ms = + ½

· Boron:

· Using quantum numbers works well for smaller atoms but can be extremely tedious for larger ones.

For example: Mg (12 electrons... not even that big an atom – can you imagine mercury with 80 electrons!!!)

n = 1 l = 0 ml = 0 ms = + ½

n = 1 l = 0 ml = 0 ms = - ½

n = 2 l = 0 ml = 0 ms = + ½

n = 2 l = 0 ml = 0 ms = - ½

n = 2 l = 1 ml = -1 ms = + ½

n = 2 l = 1 ml = -1 ms = - ½

n = 2 l = 1 ml = 0 ms = + ½

n = 2 l = 1 ml = 0 ms = - ½

n = 2 l = 1 ml = 1 ms = + ½

n = 2 l = 1 ml = 1 ms = - ½

n = 3 l = 0 ml = 0 ms = + ½

n = 3 l = 0 ml = 0 ms = - ½

So we generally use electron configurations and orbital diagrams instead:

1s2 2s22p63s2or [Ne] 3s2 or [Ne]

Energy Level Diagrams

Rules:

· aufbau principle:

· Pauli exclusion principle:

· Hund’s rule:

Energy level diagram for phosphorous, P:

The electron configuration for P is:

The orbital diagram for P is:

Exceptions to Electron Configuration Rules

*** ___________________ and ______________________ orbitals have a lower energy*** and lower energy means they are more stable and preferred. Explanation in Periodic Trends section.

Look at Chromium, Cr:

What we expect the electron configuration to be:

What the actual electron configuration is:

Why? Because this gives us ______ half-filled orbitals.

1s 2s 2p 3s 3p 4s 3d

Look at Copper, Cu:

What we expect the electron configuration to be:

What the actual electron configuration is:

Why? Because this gives us ______ half-filled orbital and _____ filled orbital.

1s 2s 2p 3s 3p 4s 3d

Notable exceptions are:s2d4 s1d5

s2d9 s1d10

Example AP MC questions:

Photoelectron Spectroscopy (PES)

Photoelectron spectroscopy is a technique used to determine the energy of electrons in an atom. It is uses a phenomenon called the photoelectric effect that we will study later in this course.

Electrons are bombarded with high-energy radiation, and the amount of energy needed to remove electrons is measured.

For this topic, you will complete the supplementary handout PES (after the AP MC questions), referring to the video link indicated.

Example AP MC questions:

Periodic Trends

1. Complete the Trends Review (in Supplemental Documents) and refer to it while studying this topic

2. Effective Nuclear Charge, Zeff

Effective nuclear charge, Zeff, is the net force of attraction between the electrons and the nucleus of the atom. It is sometimes referred to as just “nuclear charge”.

As you go across a period:

· The number of protons in each nucleus gets ____________________

· The number of electrons ____________ but the outermost electrons are in the __________________ energy level

· The shielding ________________________

· Therefore, Zeff _______________________ resulting in the outermost electrons being pulled _________ strongly

As you go down a group:

· The number of the protons in each nucleus gets ________________

· There are _________ energy levels

· The distance between the nucleus and the outermost electrons ____________________

· The shielding ____________________

· Therefore, Zeff _______________________ resulting in the outermost electrons being pulled _________ strongly

3. Important additional details concerning periodic trends:

Aufbau principle, Hund's rule and electron-pair stabilization can be used to understand a few more aspects of periodic trends. Aufbau principle indicates that in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels. Hund’s rule indicates that electrons try to stay as far apart as possible to minimize the force of repulsion between these particles. Yet, there is also stability from the pairing of oppositely spinning electrons that is also at play. In particular, there is stability is associated with symmetry (explanation beyond the scope of this course).

Consequently:

· Half-filled orbital shells – these are more stable than expected by the trend because electrons are distributed in such a way to minimize repulsions (Hund’s rule); additionally, symmetry is associated with stability

· Fully filled orbitals shells – stability associated with symmetry and often with a strong attraction of electrons to the nucleus (like in noble gases)

· Repulsion between paired electrons within a filled orbital of an asymmetrical arrangement

element

IE (kJ/mol)

Radius (Å)

P

1000

1.68

S

1012

1.83

· For example, the three electrons in the 3p orbitals on phosphorus therefore enter different orbitals with their spins aligned in the same direction. In sulfur, two electrons must occupy one of the 3p orbitals. The force of repulsion between these electrons is minimized to some extent by pairing the electrons. There is still some residual repulsion between these electrons, however, which makes it slightly easier to remove an electron from a neutral sulfur atom than we would expect from the number of protons in the nucleus of the atom.

Atomic Radius Trend

Ionization Energy Trend

Note:

Exothermic – energy (as heat) flows out of the system

Endothermic – energy (as heat) is absorbed by the system

Example AP MC questions:

If a substance is said to be:

“Diamagnetic” all electrons are paired are repelled by magnetic fields

“Paramagnetic” there are unpaired electrons are attracted to magnetic fields

Example AP question:

Valence Electrons and Ionic Compounds

Atoms with small IE tend to form ___________________, while atoms with high EA tend to form __________________.

Atoms with large differences in electronegativity tend to form ___________________ compounds when they react together.

Determining the electron configuration of an ion:

· When writing electron configurations for non-transition metal elements, remove electrons from the orbitals filled last.

Example:Na1s2 2s2 2p6 3s1 soNa+1s2 2s2 2p6 3s1 1s2 2s2 2p6

· When writing electron configurations for non-metal elements, add electrons from the orbitals filled last.

Example:Cl1s2 2s2 2p6 3s2 3p5 soCl-1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s2 3p6

· ***When writing electron configurations for transition metal elements, remove electrons from the s- orbitals first and then the d-orbitals.***

Example:Fe1s2 2s2 2p6 3s2 3p6 4s2 3d6

so

Fe2+1s2 2s2 2p6 3s2 3p6 4s2 3d6 1s2 2s2 2p6 3s2 3p6 3d6

and

Fe3+1s2 2s2 2p6 3s2 3p6 4s2 3d65 1s2 2s2 2p6 3s2 3p6 3d5

Most common ion of Fe! Why?

Example AP MC questions:

Increasing energy

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p

6d

4f

5f

Increasing energy

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p

6d

4f

5f