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Unit 2 Bonding & Atomic Structure Syllabus:

Atomic Theory • Democritus • John Dalton:  identify Dalton’s postulates and evaluate their accuracy in light of

current atomic theory. • JJ Thomson: explain how Thomson’s experiments led to his discovery of electron

properties. • Ernest Rutherford:  explain how Rutherford’s experiments led to his conclusions about

the nuclear atom. • Niels Bohr:  explain Bohr’s addition of electron energy levels to Rutherford’s nuclear

atom.  • Erwin Schrodinger •  James Chadwick

Properties of Periodic Table • explain how the contributions of scientists such as Mendeleev led to the development

of the Periodic Table as a predictive tool • explain how chemical and physical properties were historically used to develop the

Periodic Table • identify and explain properties of chemical families, including alkali metals, alkaline

earth metals, halogens, noble gases, and transition metals • use the Periodic Table to identify the chemical family of an element and explain its

properties Periodic Trends: identify and explain periodic trends, including: 

• atomic radius  • ionic radius  • electronegativity  • ionization energy • electron affinity

Lesson Description Homework

2.1 Basic Atomic Structure Guided Notes: Subatomic Particles, Rutherford Experiment, Atomic Target Practice Activity

Pg 30

2.2 Ions, Isotopes, and Weighted Atomic Mass Pg 31

2.3 General Properties of the Periodic Table Pg 31

2.4 & 2.5 Electron configuration & Quantum Numbers Pg 32-33

2.6 Energy Calculations and Flame Test Lab Pg 33

2.7 Period Trends Pg 34

2.8 Periodic Trends Activity Pg 34

2.9 Exam 2 Pg 35

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Electron Configuration • express the arrangement of all electrons in an atom using electron configuration • express the electron configuration of the d-block element • using the idea of quantum mechanics describe the areas of probability where an

electron can be located and assign the four quantum numbers • from the electron configuration identify which electrons are valence electrons and

how they are used in bonding Electromagnetic Spectrum

• describe the electromagnetic spectrum • describe the mathematical relationships between energy, frequency, and wavelength

of light • calculate wavelength and frequency using the speed of light  • calculate energy and frequency using Planck’s constant

Unit 2 Vocabulary

• quantum • quantum

mechanical model • electron

configuration • energy level • sublevel • atomic orbital • aufbau principle • Hund’s rule • Pauli exclusion

principle • Shell • Subshell • Block • Orbital notation • Electron

configuration • Bohr model • periodic law • chemical family/

group • representative

element • alkali metal • alkaline earth

metal • transition metal • inner transition

metal • halogen • Atomic radius • Ionic radius • Electronegativity • Ionization energy • Group trend • Periodic trend • Valence electrons

• electromagnetic radiation

• spectrum • wavelength () • frequency () • hertz (Hz) • Planck’s constant

(h) • speed of light (c)

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2.1 Basic Atomic Structure Guided Notes What is an Atom???

Atoms are NOT indivisible! They CAN be broken down into smaller particles. Protons, neutrons and electrons are __________evenly distributed in an atom.

• 1 amu (atomic mass unit)= 1.66 x 10-27 kg (also: 1/12th the mass of a carbon atom)

• The protons and neutrons exist in a _________________ core at the center of the atom. This is called the ___________________. The nucleus has a _______________________ charge because the protons have a positive charge and the ______________________ don’t have a charge!

• The _____________________________ are spread out around the edge of the atom. They orbit the nucleus in layers called ____________________. Electrons have a _________________________ charge.

• The ___________________ of an element contain equal numbers of protons and electrons and so have no overall charge, so if you can find it on the Periodic Table, it means it has a charge of __________________!!!

Atomic Number (Z) • The atoms of any particular element always contain the same number of _________________.

For example: hydrogen atoms always contain ____ proton and carbon atoms always contain _____ protons.

• The number of protons in an atom is known as the ______________ _______________. • It is the ______________________ of the two numbers shown in most periodic tables. • If the number of protons changes, then the atom becomes a different

____________________. • Changes in the number of particles in the nucleus (protons or neutrons) are very _________.

They only take place in ___________________ processes such as: radioactive decay, nuclear bombs or nuclear reactors.

Mass Number (A) • Electrons have a mass of almost _________________, which means that the mass of each

atom results almost entirely from the number of ______________ and neutrons in the nucleus.

• The ____________ of the protons and neutrons in an atom’s nucleus is the __________________ number. It is the _________________________of the two numbers shown in most periodic tables.

• That means in order to figure out the number of neutrons, you simply use the following:

Electrons • Atoms have no overall electrical charge

and are ____________________.

• This means atoms must have an equal number of _________________ and ____________

Subatomic Particle Charge Mass Location

Proton

Neutron

Electron

Drawing of an Atom

Mass Number= # Protons + # Neutrons

# Neutrons =Mass # - # Protons (Atomic #)

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• The number of electrons is therefore the same as the ___________________ number.

• Atomic number is the number of ____________________rather than the number of electrons, because atoms can lose or gain electrons but do not normally lose or gain protons.

2.1b Atomic Target Activity: Rutherford Scattering

Introduction: The Rutherford gold foil experiment is one of the most famous of all time. More than 25 years

after conducting the experiment, Ernest Rutherford described the results this way: “It was about as credible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you.”

The experiment itself was actually the culmination of a series of experiments, carried out over a five-year period, dealing with the scattering of high-energy alpha particles by various substances.

Ernest Rutherford received the Nobel Prize in Chemistry in 1908 for his investigation into the disintegration of the elements as a result of radioactive decay. Among the products of the radioactive decay of elements are alpha particles—small, positively charged, high-energy particles. In trying to learn more about the nature of alpha particles, Rutherford and his co-workers, Hans Geiger and Ernest Marsden, began studying what happened when a narrow beam of alpha particles was directed at a thin piece of metal foil. Alpha particles are a type of nuclear radiation, traveling at about 1/10 the speed of light. As expect4ed for such high-energy particles, most of the particles penetrated the thin metal foil and were detected on the other side. What was unexpected was that a very few of the alpha particles were actually reflected back toward the source, having been “scattered” or bent due to their encounters with the metal atoms in the foil target. The number of alpha particles that were reflected back depended on the atomic mass of the metal. Gold atoms, having the highest atomic mass of the metals studied, gave the largest amount of so-called “backscattering” Rutherford’s scattering experiments have been described as a “black box” experiment. The properties of the alpha particles, their mass, charge, speed, etc., were at least partially understood. The atoms making up the target, however, presented Rutherford with a kind of black box; the structure of the atom was not known at the time. In order to explain the results of the scattering experiment, Rutherford had to propose a new model of the atoms. A model that explained the results of the data gathered from the experiment. In 1911 Rutherford proposed the following model for the structure of the atom:

• Most of the mass of the atom is concentrated in a very small, dense central area, later called the nucleus, which is about 1/100,000 the diameter of the atom. This was proposed as a result of what data: ______________________________________________________________________________________________________________________

• The rest of the atom is apparently “empty space” This was proposed as a result of what data ______________________________________________________________________________________________________________________

• The central, dense core of the atom is positively charged, with the nuclear charge equal to about one-half the atomic mass. This was proposed as a result of what data ______________________________________________________________________________________________________________________

Objective: The purpose of this activity is to discover by indirect means the size and shape of an unknown object, which is hidden underneath the middle of a large board. By tracing the path the marble takes after

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striking the unknowntarget from a variety of angles, it should be possible to estimate the general size and shape of the unknown target.

Pre-Lab Questions

1. Read the material in your textbook about the Rutherford experiment and answer the three questions in the introduction.

2. This activity is a simulation of Rutherford’s scattering experiment. Read the entire procedure and compare the components used in this simulation to Rutherford’s original discuss what each component in our simulation corresponds to in the original experiment.

3. The key skills in this activity, as in Rutherford’s experiment, are the ability to make careful observation and to draw reasonable hypotheses. Assume that the marble strikes following sides of a possible target. Sketch the path the marble might be expected to take in each case.

! 4. Discuss what information can be inferred if the marble rolls straight through without striking the

unknown target. Materials:

1. foam board with unknown shape attached 2. marbles 3. white paper 4. push pins 5. pencil 6. ruler

Procedure: 1. Form a group of three students 2. Pin the paper to the top of the board (do not look at the shape on the underneath side) 3. Roll the marble with a moderate amount of force under one side of the board. Observe where the

marble comes out and trace the approximate path of the marble on the paper. 4. Working from all four sides of the board, continue to roll the marble under the board, making

observations and tracing the rebound path for each marble roll. Roll the marble AT LEAST 20 TIMES from each side of the box. Be sure to vary the angles at which the marble is rolled. You may use the rulers as a launching platform.

5. After sketching the apparent path from all sides and angles, the general size and shape of the unknown target should emerge.

6. Form a working hypothesis concerning the structure of the unknown target. Based on this hypothesis, repeat as many “targeted’ marble rolls as necessary to confirm or revise the structure.

7. Check your answer with your teacher. DO NOT look under the board. 8. If time permits try an extension, or another shape.

Post Lab Questions 1. Draw the general size and shape of the target to approximate scale in the square below.

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2. The speed of the marble rolls was an uncontrolled variable in this activity. How would the outcome of the scattering test have been different if the marble speed had been faster or slower?

3. Compare the overall size of the target with the size of the marble used to probe its structure. How would the outcome of the scattering test have been different if different size marbles had been used? Explain.

2.2 Ions, Isotopes, and Weighted Atomic Mass Isotopes A) Atoms with same number of protons ( #p+ ) but different numbers of neutrons (#no). Isotopes have the same atomic number (Z) but different mass numbers.

B) Recognizing Isotope Notations i) X = element from periodic table

ii) A = mass number = #p + #n iii) Z = atomic number = #p -NOTE: In a neutral atom, Z = #p = #e

C) Examples Sodium-23 Na-23 (Z=11) 11 p, 12 n, 11 e

Sodium 24 Na-24 (Z= 11) 11 p, 13 n, 11 e

Ions A) Ions are charged atoms. B) Charges arise due to loss or gain of electrons. C) If atom loses electrons, a positively charged cation is formed.

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Example: Li > Li+ + 1 e-

D) If atom gains electrons, a negatively charged anion is formed.

Example: F + 1 e- > F- E) In ordinary matter, cations and anions always occur together so that matter is charge-neutral overall.

Atomic Mass 1) Atomic Mass of Element represents the weighted average of all the naturally occurring isotopes of that element. 2) Value found below the atomic symbol of the element in the periodic table.

A) Calculating Atomic Mass for an Element

i) Must be given relative abundance (%) for each isotope ii) Must know isotopic mass for each isotope iii) Convert percent abundance to a decimal (divide by 100%) iv) Use following formula to obtain result

Atomic Mass = Σ (fraction of isotope n) x (mass isotope n)

Example: Find Atomic Mass of Carbon given the following data.

Data: 98.89% Carbon 12 isotopic mass 12 amu 1.11% Carbon 13 isotopic mass 13.0034 amu

Solution: (12 amu * 0.9889) + (13.0034 amu * 0.0111) = 11.867 amu + 0.144 amu = 12.011 amu

Atomic Weights/Mass

Naturally occurring magnesium consists of three stable isotopes:

What is the atomic weight of Magnesium? Ans:_______________________

Naturally occurring silicon consists of three stable isotopes:

isotope amu Abundance

Mg-24 23.985 78.99%

Mg-25 24.986 10.00%

Mg-26 25.983 11.01%

isotope amu Abundance

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What is the atomic mass of Si-30? Ans:_______________________

Protons, Neutrons, Electrons Fill in the table below with the correct numbers (first one is done as an example)

2.3 General Properties of the Periodic Table Mendeleev's Periodic Table (1869) A. Organization

1. Vertical columns in atomic weight order a. Mendeleev placed elements in rows with similar properties

2. Horizontal rows have similar chemical properties B. Missing Elements

B . Gaps existed in Mendeleev’s table a. Mendeleev predicted the properties of the “yet to be discovered” elements

Si-28 27.977 92.21%

Si-29 28.976 4.70%

Si-30 ? 3.09%

Symbol name Atomic number

Mass number

charge # of parts in nucleus

# of protons

# of neutrons

# of electrons

12 25 0

1- 35 18

7 7 10

K4119

Na2311

+141

19K

K3919

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(1) Scandium, germanium and gallium agreed with his predictions C. Unanswered Questions

1. Why didn't some elements fit in order of increasing atomic mass? 2. Why did elements exhibit periodic behavior?

Moseley and the Modern Periodic Table (1911) A. Protons and Atomic Number

1. The periodic table was found to be in atomic number order, not atomic mass order B. The Periodic Law

1. The physical and chemical properties of the elements are periodic functions of their atomic numbers

2. Elements with similar properties are found at regular intervals within the periodic table * Moseley was killed in battle in 1915, during WWI. He was only 28 years old

Organization of the Table 1. Groups or Families

a. Vertical columns containing elements with similar chemical properties 2. Periods (series)

a. Horizontal rows of elements 3. Metals and Nonmetals

a. A stair-step line on the table separates the metals from the nonmetals b. Metalloids (Semimetals) straddle the line and have properties of both metals

and nonmetals 4. Lanthanide and Actinide Series (Inner Transition Metals)

a. Metals and man-made metal elements 5. Group 1 – Alkali metals (the most reactive metal elements) (except hydrogen (H)

also in this group) 6. Group 2 – Alkaline earth metals (very reactive metal elements) 7. Group 17 – Halogens (the most reactive nonmetal elements) 8. Group 18 – Noble gases (the least reactive elements – inert and very stable)

Types of Elements A. Metals

1. Luster 2. Good conductors of heat and electricity 3. Malleable 4. Ductile 5. High tensile strength

B. Nonmetals 1. Many nonmetals are gases at room temperature 2. Solid nonmetals tend to be brittle and non-lustrous 3. Poor conductors of heat and electricity

C. Metalloids 1. Some properties of metals and some properties of nonmetals 2. Solids at room temperature 3. Semiconductors of electricity

D. Noble Gases 1. All are gaseous members of group 18

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2. Generally unreactive and stable

Periodic Table

• Prior to 1860 no agreement/method to accurately determine masses of atoms.

• First International Congress of Chemists –1860– Stanislao Cannizzaro presented method for

accurately measuring atomic masses– Looked for relationships between atomic

masses and other properties of elements

John Newlands

• Noticed elements properties repeated every 8th element when arranged by atomic mass

• Named this phenomenon “the Law of Octaves”

• Did not work for all elements

• First tables arranged elements by atomic weight – Could not agree on atomic weights therefore

tables were different

Julius Lothar Meyer

• Developed first modern table– Consisted of 28 elements

divided into 6 families– Families (groups) had similar

chemical and physical properties

– Discovered all elements in same family had same number of valence e- --outermost electrons in highest energy level

– Why?

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• Why can most elements be arranged by atomic mass?

• What was the reason for chemical periodicity?

Dmitri Mendeleev• Noticed that properties

repeat themselves at certain intervals

• Arranged all knownelements into one table based on properties– 1869

• 1871 - Proposed the “Periodic Law”

• Based on the properties spaces were left for unknown elements (Sc, Ga, Ge)

• Upon discovery of other elements inconsistencies were found with Mendeleev’s table

• Atomic masses improved and they no longer arranged the elements by increasing atomic mass

Henry Mosely

• Discovered elements contain unique number of protons (atomic number) - 1911

• Arranged elements by atomic number -1913

• Fully explained the Periodic Law

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2.4 Quantum Numbers & Electron Configuration

Quantum Numbers and Electron Configuration

1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A 8A

Quantum Mechanical Model• Electrons are attracted to the nucleus by

electrostatic forces between oppositely charged objects

• Electrons reside in space that are different distances from nucleus

• Limited number of regions where an e- can reside (energy is quantized)

• Atoms absorb/emit radiation when e- move

Principal Quantum Number (n)• Indicates the relative sizes and energies of

atomic orbitals• n is an integer greater than 1 (n =1, 2, 3,….)• 7 energy levels have been identified for

Hydrogen

1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A 8A

Continued…• Energy states have negative values• Energy values increase (become more

positive) farther from nucleus n= ∞ has 0 energy value

• Can completely remove an electron from an atom when n=∞ (an ion is formed)

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1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A 8A

s

p

d

f

Magnetic Quantum Number (ml)• Describes the orientation of an orbital around

the nucleus• s orbital is spherical and has only 1 orientation

(m = 0)• p orbital can orient along each axis (x, y, &

z)(m = -1, 0, +1)

d - subshell• d – subshell contains 5 orbitals

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1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A 8A

Spin Quantum # (ms)• spin makes the electron behave like a tiny

magnet • spin can be clockwise or counterclockwise • spin quantum number can have values of +1/2

or -1/2

• Can only put 2 electrons in each orbital and must have opposite spins

Predicting Electron Configuration• Atoms like to have the most stable configuration as

possible.• Number of subshells equal to shell number, n• In order of increasing energy subshells labeled s, p, d,

& f4s < 4p < 4d < 4f

• Always odd number of orbitals in a subshell• Maximum number of e- in subshell equal 2x number

of orbitals• Electrons are added to an atom, one at a time,

starting with lowest available orbital – Aufbau principle

Pauli Exclusion Principle• No two electrons in the same atom can have

the same set of four quantum numbers. • Electrons must have opposite spins in the

same orbital• Spins of electrons represented with

1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A 8A

1s

2s2

p3s 3p

4s3d 4p

5s 4d 5p

6s 5d 6p

7s 6d 7p

4f

5f

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Orbital Notation• Unoccupied orbital is designated with a

_____ and the orbitals name written underneath

3p 3p 3p• Electrons are placed in each orbital using• Give the orbital notation for carbon and

fluorine

Electron Configuration Notation• Number of electrons in a sublevel is shown by

adding a superscript to the sublevel designation

• Boron has 5 electrons– 1s22s22p1

• Give the proper configurations for carbon and fluorine

Carbon• 6 electrons

1s 2p2s

fluorine• 9 electrons

1s 2p2s

Summary video on electron configuration

• http://www.youtube.com/watch?v=Vb6kAxwSWgU

Short Hand Notation• Locate the Noble gas preceding your element– The noble gas has a full outer shell in its electron

configuration• Place the symbol for the Noble gas in brackets– [X]

• Complete the electron notation for the desired element

Example• Calcium

• Noble Gas preceding Ca is Argon (Ar)• [Ar]4s2

• Give short hand notation for Bromine• [Ar]4s23d104p5

1s 2p2s 3s 4s3p

• Carbon: 1s22s22p2

• Fluorine: 1s22s22p5

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Electron Configuration Practice Worksheet In the space below, write the unabbreviated electron configurations for the following elements:

1) Sodium ________________________________________________

2) Iron ________________________________________________

3) Bromine ________________________________________________

In the space below, write the abbreviated electron configurations for the following elements:

6) Cobalt ________________________________________________

7) Silver ________________________________________________

8) Tellurium ________________________________________________

Determine what elements are denoted by the following electron configurations:

11) 1s22s22p63s23p4 ____________________

12) 1s22s22p63s23p64s23d104p65s1 ____________________

13) [Kr] 5s24d105p3 ____________________

Determine which of the following electron configurations are not valid:

14) 1s22s22p63s23p64s24d104p5 ____________________

15) 1s22s22p63s33d5 ____________________

16) [Ra] 7s25f8 ____________________

2.5 Energy Calculations & Flame Test Lab

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Wave Description of Light• Electromagnetic radiation (ER): form of

energy that exhibits wavelike behavior as it travels through space

• All forms of ER together make the electromagnetic spectrum

• All forms of ER move at a constant speed of about 3.0 x108 m/s (speed of light, c)

• Wavelength (λ): distance between corresponding points on adjacent waves.

• Frequency (ν): number of waves that pass a specific point in a given time, usually one second.

• Unit: Hertz (Hz), aka (1/s) or (s-1)• For electromagnetic radiation, frequency and

wavelength are relatedC=λν

• If λ increases, what must happen to ν? Does cchange?

Particle description of light• Max Planck, 1900s, suggested that object emit

energy in small, specific amounts called quanta

• Quantum: minimum amount of energy that can be gained or lost by an atom

• 1905 Einstein expands on this idea. ER have dual wave/particle nature.

• While light emits many wavelike particles, it can also be thought of as a stream of those particles

• Einstein named the particles photons• Photon: particle of ER having zero mass and

carrying a quantum of energy

• Planck proposed a relationship between a quantum of energy and the frequency of radiation

E=hν

• E is energy, in Joules, of a quantum of radiation

• h is Planck’s constant (fundamental physical constant)= 6.626 x10-34 J*s

• ν is freqency of radiation

• Energy of a particular photon depends on the frequency of the radiation

Ephoton= hν• Einstein’s Explanation: ER is absorbed by matter

only in whole numbers of photons– For e- to be ejected, must be struck by single photon

possessing at least minumum energy– According to equation, this energy corresponds to

frequency– If photon’s frequency is below minimum, no e- ejected

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Energy Calculations Practice:

The Hydrogen Atom• Ground state: lowest energy state of an

atom• Excited state: atom has higher potential

energy than ground state• When an excited atom returns to ground

state, it gives off energy it gained in the form of ER.

• Production of colored signs (neon) is example

Experiment• Pass electrical current through H gas in

vacuum tube at low pressure• Emits characteristic pink glow• When the light was passed through prism,

separated into series of specific frequencies and therefore wavelength (… equation?) of visible light

• These bands are Hydrogen’s line emission spectrum

Problem?• Classical theory predicted that H atoms would

be excited by whatever amount of energy was added to them.

• Expected to observe continuous range of frequencies of ER, or continuous spectrum.

• Why had H only given off specific frequencies of light?

• ! Quantum Theory

• Excited H atom falls back from excited state to ground state, and emits photon

• This energy is equal to difference between initial and final state

Implications• Since H atoms emit only specific frequencies,

difference between energy states must be fixed.

• Therefore, e- of H atom exists only in very specific energy states

• In 1913, Bohr proposed a model that linked the atom’s electron with photon emission

• Energy is higher in orbits farther from nucleus (like a ladder)

• Based on the wavelengths of hydrogen’s line emission spectrum, Bohr calculated energies the e- would have in the allowed energy level for H atom.

• Bohr’s calculated values agreed with experimentally observed values for lines in each series

• Scientists tried to apply this model to other element’s atoms

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1. What is the energy of a mole of photons with a wavelength of 1.60 10-3 m?

2. What is the energy of a photon with a wavelength of 2.65 10-4 m?

3. What is the frequency of a photon with a wavelength of 3.70 10-6 m?

4. What is the frequency of a photon with a wavelength of 8.60 103 nm?

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2.6b FLAME TEST LAB

In this activity, you will investigate the colors of flame produced by solutions of metal salts. When a substance is heated in a flame, the atoms absorb energy from the flame. This absorbed energy allows the electrons to be promoted to excited energy levels. From these excited energy levels, there is a natural tendency for the electrons to make a transition or drop back down to the ground state. When an electron makes a transition from a higher energy level to a lower energy level, a particle of light called a photon is emitted. Both the absorption and emission of energy are quantized – only exact amounts of energy are required.

An electron may drop all the way back down to the ground state in a single step, emitting a photon in the process. Alternatively, an electron may drop back down to the ground state in a series of smaller steps, emitting a photon with each step. In either case, the energy of each emitted photon is equal to the difference in energy between the excited state and the state to which the electron relaxes. The energy of the emitted photon determines the color of light observed in the flame.

In this activity, metal salts in alcohol are burned, producing different colored flames. By comparing the color given off by an unknown with the known metal salts, the identity of the metal salt can be determined.

Flame Tests Activity Materials:

• matches • Spray bottle(s) from your teacher • ethanol • the following metal salts • lithium chloride • strontium chloride

• calcium chloride • copper(II)chloride • sodium chloride • potassium chloride • UNKNOW

Procedure: 1. Obtain a spray bottle of each metal ion. 2. Light your Bunsen burner. 3. If your flame blows out TURN OFF the gas immediately and relight the burner. 4. Holding the bottle upright and approximately 12 inches from the flame squirt the contents

at the flame. It may be necessary to move the location and the distance of the squirting depending on air currents in the room.

5. Record the color of the resulting flame and the intensity, be descriptive in your colors. 6. Repeat for all the salts.

Data Table:

Metal found in the salt

Flame Color and Intensity

Lithium

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Flame Test Analysis 1. List the colors observed in this lab from the highest energy to the lowest energy.

2. List the colors observed in this lab from the highest frequency to the lowest frequency.

3. List the colors observed in this lab from the shortest wavelength to the longest

wavelength.

Strontium

Calcium

Copper

Sodium

Potassium

UNKNOWN

Identity:_____________

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4. What is the relationship between energy, frequency, and wavelength?

5. Do you think we can use the flame test to determine the identity of unknowns in a mixture? Why or why not?

6. How are electrons “excited” in this part of the experiment? What does it mean when the electrons are “excited”?

7. Explain why we did not see distinct lines (like on an emission spectrum) when the metal salts were burned.

8. Why do different chemicals emit different colors of light?

9. Why do you think the chemicals have to be heated in the flame first before the colored light is emitted?

10. Colorful light emissions are applicable to everyday life. Where else have you observed colorful light emissions? Are these light emission applications related? Explain.

2.7 General Period Trends

Trends in the Periodic Table

Periodic Law

• The physical and chemical properties of the elements are periodic functions of their atomic numbers.

• Aka – when elements are arranged by increasing atomic number, elements with similar properties appear at regular intervals.

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s-Block Elements

• Groups 1 & 2• All elements in group 1 & 2 will have an

electron configuration of – ns1 or ns2 where n = highest energy level

occupied

Alkali Metals

• Group 1 elements• In the elemental state

– Soft– Silvery metal– High melting points– Extremely reactive therefore are not found in

elemental state in nature• React violently with water to produce

hydrogen gas

Parts….

• Alkali metals – group 1• Alkaline earth metals – group 2• Halogens – group 17• Transition metals – d block elements• Inner Transition metals

– Lanthanides (elements 58-71) added in early 1900’s

• Have very similar properties– Actinides (elements 90-103)

Hydrogen & Helium

• H has same valence electrons as group 1 but does not share any other properties

• He share same electron configuration (valence e-) as group 2 but does not share same properties– Placed with group 18 because it is very stable

Alkaline – Earth Metals

• Group 2 elements• Outer most s orbital is full

– Do not exhibit stability (outer p orbital is empty)

• Properties– Harder, denser than group 1– Higher melting points than group 1– Not as reactive but too reactive to be found in

nature in elemental form

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p-block elements

• All elements in p block have a full s orbital• Properties

– Contain all non metals except H & He– Contain all metalloids (exhibit properties of both

metals and non metals)• Have semi conducting properties

– Contains 6 metals

• Elements in s & p block make up the representative elements

Halogens

• Group 7A/17– Most reactive non metals (Fluorine is most

reactive)– Will bond with a metal to form a salt– F & Cl are gases at room temp– Br is a liquid at room temp– I & At are solids at room temp

Exceptions in the d-block

• The following elements have odd configurations– Cr: [Ar]4s13d5

– Cu: [Ar]4s13d10

– Ag: [Kr]5s14d10

• More stable with half filled s & d orbitals or full d orbital

• Exceptions follow throughout the d element similar to Chromium and Copper

d-block elements

• Transition elements – Beginning filling the 3d orbitals– Good conductors of electricity– High luster– Less reactive than s-block elements

• Can be found in elemental form

Trends in Atomic Size cont.• Group - atomic radius increases as you go down a

group.

Why? • There is a significant jump in the size of the nucleus

(protons + neutrons) each time you move from period to period down a group.

• Additionally, new energy levels of elections clouds are added to the atom as you move from period to period down a group, making the each atom significantly more massive, both is mass and volume.

Atomic Radius• Atomic radius is simply the radius of the atom,

an indication of the atom's volume.• Atomic radius is one-half the distance between

the two nuclei in a molecule consisting of two identical atoms.

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Electronegativity Trends• Period - electronegativity increases as you go from left

to right across a period.

• Why? Elements on the left of the period table have 1 -2 valence electrons and would rather give those few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's electrons. As a result, they have low electronegativity. Elements on the right side of the period table only need a few electrons to complete the octet, so they have strong desire to grab another atom's electrons.

Electronegativity• Electronegativity is an

atom's 'desire' to grab another atom's electrons.

Ionization Energy

• Ionization energy is the amount of energy required to remove the outermost electron/s.

• Ionization energy is closely related to electronegativity.

Electronegativity Trends cont.• Group - electronegativity decreases as you go down a

group.

• Why? Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have so many electrons that loosing or acquiring an electron is not as big a deal. This is due to the shielding effect where electrons in lower energy levels shield the positive charge of the nucleus from outer electrons resulting in those outer electrons not being as tightly bound to the atom.

1) How easily electrons can be removed (ionization energy)

from an atom

Ionization Energy Trends cont.

• Group - ionization energy decreases as you go down a group.

• Why? The shielding effect makes it easier to remove the outer most electrons from those atoms that have many electrons (those near the bottom of the chart).

Reactivity

Reactivity refers to how likely orvigorously an atom is to react with other substances. This is usually determined by two things:

Ionization Energy Trends

• Period - ionization energy increases as you go from left to right across a period.

• Why? Elements on the right of the chart want to take others atom's electron (not given them up) because they are close to achieving the octet. The means it will require more energy to remove the outer most electron. Elements on the left of the chart would prefer to give up their electrons so it is easy to remove them, requiring less energy (low ionization energy).

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Period - reactivity decreases as you go from left to right across a period.

Group - reactivity increases as you go down a group

Why? The farther to the left and down the periodic chart you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity.

Reactivity of Metals

The transfer/interaction of electrons is the basis of chemical

reactions.

Period - reactivity increases as you go from the left to the right across a period. Group - reactivity decreases as you go down the group.

Why? The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electron.

Reactivity of Non-Metals

2) or how badly an atom wants to take other atom's electrons

(electronegativity)

Ionic Radius vs. Atomic Radius

• Metals - the atomic radius of a metal is generally larger than the ionic radius of the same element.

• Why? Generally, metals loose electrons to achieve the octet. This creates a larger positive charge in the nucleus than the negative charge in the electron cloud, causing the electron cloud to be drawn a little closer to the nucleus as an ion.

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2.8 Periodic Trends Activity Determining Periodic Trends Activity

Purpose To understand the periodic nature of the periodic table – by knowing an element’s location on the periodic table one can determine the element’s number of valence electrons and have an awareness of its ionization energy, density, electron affinity, electronegativity, atomic radius, and ionic radius, relative to other elements

Materials Tiles from Periodic Trends, periodic table, scissors, tape, white paper, graph paper

Procedure Part 1

1. You will be given cards for each of five elements in a certain group. Each card contains information about the physical properties of the element. Use the descriptions in the Periodic Trends section below to place your element cards in what you think is the correct arrangement. You only have representative or main-group elements. The transition metals and noble gases are not used in this activity.

2. Use a periodic table to predict the atomic number and period of each element and write them in the spaces provided on the card.

Part 2 1. As a group, organize all the elements into a periodic table. Tape the elements onto

white paper when you are sure of their placement. 2. Compare your periodic table with the teacher’s key and make any necessary

adjustments. 3. Complete the Analysis and Conclusions section.

Terminology of Periodic Trends Valence electron: An electron in an outer shell of an atom that can participate in forming chemical bonds with other atoms. Nonvalence electrons are tightly bound to the nucleus and are called core electrons. Octet rule: Atoms tend to gain, lose, or share electrons to reach eight electrons in their outer electron shells. Ionization energy: The amount of energy required to remove one electron from a neutral atom that is in the gaseous state. Elements on the right side really want to keep their electrons because they are so close to achieving an octet. Elements on the left side don’t mind losing an electron. Going down a group, the shielding effect of additional electrons makes it easier to remove outermost electrons. Atomic radius: One-half the distance of a single bond between two atoms of an element. As we go down a group, the principal energy level increases, causing an increase in the average distance between the electrons and the nucleus. Atoms get bigger as electrons are added to the principal energy levels. All the atoms in a given period have their outermost electrons in

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the same principal energy level. As we go across a period, we are adding electrons to the same principal energy level. The atoms do not get bigger across a period because we are also adding protons to the nucleus as we move across a period. The additional protons in the positive nucleus increase the pull on the negative electrons resulting in smaller atoms across a period. Ionic radius: One-half the center-to-center distance between two ions in a crystalline ionic compound. Ions on the left side of the table are smaller than the atoms they come from because they’ve lost an electron. Ions on the right side are bigger than the atoms they come from because they’ve gained an electron. Electronegativity: A measure of attraction of an atom for a pair of shared electrons, a measure of whether an atom “grabs” more than its fair share of the shared electrons. As you go down a group, atoms have more total electrons so they don’t really care that much about their outermost ones. As you go across a period, atoms have more valence electrons and are so close to achieving an octet that they’ll grab another atom’s electrons.

Analysis and Conclusions 1. Use the words increases or decreases to describe the trends for the following properties of the representative elements, going from left to right across a period:

• ionization energy ____________________________________

• atomic radius ____________________________________

• ionic radius ____________________________________

• electronegativity ____________________________________

2. Use the words increases or decreases to describe the trends for the following properties of the representative elements, going from top to bottom down a group:

• ionization energy ____________________________________

• atomic radius ____________________________________

• ionic radius ____________________________________

• electronegativity ____________________________________

3. When elements in Group I ionize, are they more likely to gain or lose electrons? Explain your answer.

4. Explain why Group I elements have ionic radii smaller than their atomic radii.

5. When elements in Group VII ionize, are they more likely to gain or lose electrons? Explain your answer.

6. Explain why Group VII elements have ionic radii larger than their atomic radii.

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7. Compare how the melting point for metals and nonmetals generally changes as you move down a group.

8. What is the relationship between number of valence electrons and group number?

9. Create a graph of ionization energy versus atomic number and identify the different periods. Compare the ionization energies of metals to nonmetals.

10. Create a graph of atomic radius versus atomic number and identify the different periods. What is the trend going down a group? What is the trend as you go across a period?

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Unit 2: Homework

Hwk 2.1 Atomic Structure

Rutherford’s Gold Foil Experiment: view experiment in video online

Summarize the experiment and the major conclusions below: ____________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________

Atomic theory, as we know it today, is the result of the contributions of many scientists who did the research, disproved old models and suggested the new ones, and added to the atomic theory. The items below are the concepts that were developed in the history of the atomic theory. Rank them in order from oldest to most recent.

____Plum Pudding Model ____Electrons are in orbits ____Atoms of the same element have the same properties ____Positivity-charged atomic nucleus

Indicate whether each of the following statements about the nucleus of an atom is true or false. _____The nucleus of an atom is neutral.

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_____The nucleus of an atom contains only neutrons. _____The number of nucleons present in the nucleus is equal to the number of electrons present outside the nucleus. ______The nucleus accounts for almost all the volume of an atom. ______The nucleus accounts for almost all the mass of an atom. ______The nucleus can be positively or negatively charged depending on the identity of the atom.

Hwk 2.2 Ions, Isotopes and average atomic weight

Complete this table. Note that the atoms/ions are not necessarily neutral.

Calculate the atomic mass of naturally occurring tungsten (W) to five sigfigs given the following

isotopic masses and abundances: W = ______________________ amu

Suppose that a fictitious element, X, have two isotopes: (59.015 amu) and (62.011 amu). The lighter isotope has an abundance of 73.7%. Calculate the atomic mass of the element X.

Hwk 2.3 General Properties of Periodic Table

Identify these elements based on their location in the periodic table. Give the symbol and number Period 5, group 14 (4A) __________________________________________________

Element Se

Mass Number 76 65

Number of Neutrons 36 39

Number of Protons 36

Number of electrons 36

Charge -2 +1

Isotope

Isotopic mass (amu)

Abundance )%)

179.946706 0.12

181.948206 26.50

182.9502245 14.31

183.9509326 30.64

185.954362 28.43

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Period 4, group 12 (2B)___________________________________________________ Period 5, group 18 (8A) ___________________________________________________

Label these groups on the periodic table below: Halogens, Alkali metals, Alkaline Earth metals, Noble Gases, Transition Metals, Hydrogen, Helium, Metalloids, Non-metals, Lanthanide and Actinides. You can use arrows of shade the blocks in color.

Classify each of the following elements as Metals (M), Metalloids (MO) and Non-metals (NM) As, _______ Ga________, Tl_________, Xe_______, Si_______, S________, Bi_________

Hwk 2.4 Electron Conf & Quantum Numbers Identify the atom with the following ground-state electron configuration for its valence shell. Symbol: ____________

Classify the following orbitals as s, p, d or f, according to their shapes:

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What quantum numbers specify these subshells?

Hwk 2.5 More Electron Conf & QN

In the space below, write the unabbreviated electron configurations for the following elements:

1) Barium ________________________________________________

2) Neptunium ________________________________________________

In the space below, write the abbreviated electron configurations for the following elements:

1) Radium ________________________________________________

2) Lawrencium ________________________________________________

Determine what elements are denoted by the following electron configurations:

1. [Xe] 6s24f145d6 ____________________

2. [Rn] 7s25f11 ____________________

Determine which of the following electron configurations are not valid:

1. [Kr] 5s24d105p5 ____________________

1. [Xe] ____________________

Hwk 2.6 Energy Calculations

1. What is the energy of a photon with a wavelength of 6.55 102 nm?

2. What is the wavelength of a photon with an energy of 3.40 10-18 J?

4s 2p 3d 4f

n =

l =

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3. What is the wavelength of a photon with a frequency of 2.35 1012 s

Hwk 2.7 Periodic Trends

Rewrite these elements from most to least electronegative. Al, Na, Rb, F, N ______________________ Rank these elements from largest to smallest radius. Ne, Li, B, N, F, C, O, Be _______________________ Rank these elements from highest to lowest ionization energy. Br, Kr, K, Ge, Ca, Se _________________ Rank these ions according to largest to smallest radius. N-3, F-, Mg2+,O2-, Na+ _______________________

Hwk 2.8 More Periodic Trends

1. What is the symbol for the following elements. a. Magnesium _____________ b. Potassium ______________

2. What are the names of the following elements. a. C __________________ b. Cl _________________

3. What period are the following elements in? a. He _______________ b. Ge _________________

4. What group are the following elements? a. Sulfur _______________ b. Ca _________________

5. Give me an atom with the following characteristics. a. Halogen _________________ b. Nonmetal ________________

c. Alkali metal ______________ d. metalloid ________________

e. Lanthanide series __________ f. Alkaline Earth metal ________________

g. Transition metal ___________ h. Nobel gas ________________

6. Write the electron longhand and shorthand configurations and give the quantum #s for: a. Li _______________________________________________________________

b. Na ______________________________________________________________

c. K _______________________________________________________________

7. What are valence electrons? __________________________________________________

8. How many valence electrons are in the following element?

a. F ________ b. Cl ___________ c. Br ____________ d. I _____________

e. O ________ f. S ___________ g. Se ____________ h. Te ____________

Hwk 2.9 Review 1. When compared to copper-63, copper-65 has more —

a. protons b. neutrons c. energy levels d. bonding configurations

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2. Which of the following pairs of atomic symbols represent isotopes of the same element? a. 235U and 238U b. P4 and P8 c. 32P and 83Pb d. 50Sn

and 51Sb

3. Which of the following represents a particle containing 8 protons, 9 neutrons and 8 electrons?

A. Oxygen-16 B. Nitrogen-17 C. Oxygen-17 D. Fluorine-15 4. Fill in the chart below: Remember:

5. Define a family. _______________________________________________________

6. What is a period? ________________________________________________________

7. What is the symbol for the following elements. a. Magnesium _____________ b. Potassium ______________

8. What are the names of the following elements. a. C __________________ b. Cl _________________

9. What period are the following elements in? a. He _______________ b. Ge _________________

10. What group are the following elements? a. Sulfur _______________ b. Ca _________________

11. Give me an atom with the following characteristics. a. Halogen _________________ b. Nonmetal ________________

c. Alkali metal ______________ d. metalloid ________________

e. Lanthanide series __________ f. Alkaline Earth metal ________________

g. Transition metal ___________ h. Nobel gas ________________

12. Write the electron longhand and shorthand configurations and give the quantum #s for: a. Li _______________________________________________________________

b. Na ______________________________________________________________

c. K _______________________________________________________________

13. What are valence electrons? __________________________________________________

14. How many valence electrons are in the following element?

a. F ________ b. Cl ___________ c. Br ____________ d. I _____________

Isotope Symbol

Isotope name

Mass number

#of particles in the nucleus

Number of protons

Number of neutrons

Number of electrons

Sodium-23 23 23

32

8 10

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e. O ________ f. S ___________ g. Se ____________ h. Te ____________

On the blank periodic table below 15. Label the s, p, d, and f block elements 16. Create a circle that fills the whole box where the largest atom exists in the periodic

table. 17. Put a dot where the smallest atom is in the periodic table. 18. Put a triangle on the box with the atom with the highest electronegativity 19. Put a square in the box with the lowest ionization energy 20. Label with arrows the trends for: atomic radius, ionic radius (metals and nonmetals),

ionization energy, and electronegativity

From Electron Configuration Notes:

17. What is the shape of the s orbital? _________p orbital?___________ d orbital? _________________

18. Which of the following orbitals is closest to the nucleus? a. 2s b. 3p c. 1s d. 4d

19. In the wave-mechanical (quantum) model of the atom, orbitals are regions of the most probable locations of:

a. protons b. positrons c. neutrons d. electrons

20. (CHALLENGE) Heiseinberg’s Uncertainty Principle states: ________________________________________________________________________________________________________________________________________________________

21. Identify the following atom, 1s22s22p63s23p64s23d6 _________________________

22. Write out the orbital notation and give the quantum numbers for F, N, Zn:

23. Write the electron configuration for the above elements:

24.Write the shorthand (noble gas) notation for the above elements:

F: __________________________________________

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F1— : __________________________________________

Zn: __________________________________________

Zn2--: _________________________________________

Sr: _________________________________________

Sr1+: _________________________________________