Learning Objectives Atomic Theory and Periodic Table

Periodic Table and Electrons

Chemistry

Essential knowledge and skills:

Distinguish between a group and a period.
Identify key groups, periods, and regions of elements on the periodic table.
Identify and explain trends in the periodic table as they relate to ionization energy, electronegativity, shielding effect, and relative sizes.
Compare an element’s reactivity to the reactivity of other elements in the table.
Relate the position of an element on the periodic table to its electron configuration.
Determine the number of valence electrons and possible oxidation numbers from an element’s electron configuration.
Write the electron configuration for the first 20 elements of the periodic table.

Essential understandings:

The periodic table is arranged in order of increasing atomic numbers.
The names of groups and periods on the periodic chart are alkali metals, alkaline earth metals, transition metals, halogens, and noble gases.
Metalloids have properties of metals and nonmetals. They are located between metals and nonmetals on the periodic table. Some are used in semiconductors.
Periods and groups are named by numbering columns and rows. Horizontal rows called periods have predictable properties based on an increasing number of electrons in the outer energy levels. Vertical columns called groups or families have similar properties because of their similar valence electron configurations.
The Periodic Law states that when elements are arranged in order of increasing atomic numbers, their physical and chemical properties show a periodic pattern.
Periodicity is regularly repeating patterns or trends in the chemical and physical properties of the elements arranged in the periodic table.
Atomic radius is the measure of the distance between radii of two identical atoms of an element. Atomic radius decreases from left to right and increases from top to bottom within given groups.
Electronegativity is the measure of the attraction of an atom for electrons in a bond. Electronegativity increases from left to right within a period and decreases from top to bottom within a group.
Shielding effect is constant within a given period and increases within given groups from top to bottom.
Ionization energy is the energy required to remove the most loosely held electron from a neutral atom. Ionization energies generally increase from left to right and decrease from top to bottom of a given group.
Electron configuration is the arrangement of electrons around the nucleus of an atom based on their energy level.
Electrons are added one at a time to the lowest energy levels first (Aufbau Principle). Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results (Hund’s Rule).
Energy levels are designated 1–7. Orbitals are designated s, p, d, and f according to their shapes and relate to the regions of the Periodic Table.
An orbital can hold a maximum of two electrons (Pauli Exclusion Principle).
Atoms can gain, lose, or share electrons within the outer energy level.
Loss of electrons from neutral atoms results in the formation of an ion with a positive charge (cation). Gain of electrons by a neutral atom results in the formation of an ion with a negative charge (anion).

ELEMENTS AND THE PERIODIC TABLE
WHAT’S IN THE NAME?

Provide the name and chemical symbol for the element, which sounds similar to a possible correct answer. See number one for an example.

CLUEELEMENTCLUE ELEMENT
1. well drillers decisionboron B24. a spice

2. to press laundry25. a blitz by police

3. policeman26. …..on the Range

4. mother’s sister’s money 27. …of Arabia

5. where dishes are washed28. dull chemistry lecture

6. a foolish prisoner29. Golden Gate Bridge state

7. natives of North America30. …bladder or …stones

8. water and gin31. European country

9. shown the way32. another European country

10. I sit down to eat33. repair clothes

11. a Ford product34. 50 per cent

12. a popular house plant35. larger than a coyote

13. have the sniffles36. God of the underworld

14. United States citizen37. God of the sea

15. don’t take any wooden38. a good . . . (helper)

16. playing a part or role39. to brown a roast

17...... pop40. girl’s names calcium Ca

18. a lisper saying “listening” (many options!)

19. technician

20. Lone Ranger’s horse

21. endure or tolerate pain

22. what you walk on at home

23. leg joint above calf

CLUE: ium ending read as “him”ELEMENT

1. doctors do thishelium (heel him)

2. doctors do this too

3. funeral homes do this

4. cowboys do this to horses

5. “Anything to keep him quiet” mother
says to father.

6. grab him

7. to get him off my back

History of the Periodic Table

J.A.R. Newlands - 1867 first version of Periodic Table. Newlands arranged the known elements by increasing atomic mass along horizontal rows seven elements long, stated that the 8th element would have similar properties to the first from the series. Newlands called this the law of octaves. Newlands' work failed after Ca in predicting a consistent trend.

Dimitri Mendeleev 1869, Professor of Chemistry at the University of Saint Petersburg (Leningrad). Mendeleev stated that the elements vary periodically (in cycles) according to their atomic masses.

Mendeleev separated his elements and left spaces on his table in order for the periodicity to continue. He then predicted that elements would be discovered to fill these "gaps" in the table. Mendeleev even accurately stated the properties of these elements. Scandium(eka-boron), gallium(eka-aluminum), and germanium(eka-silicon). By 1886 all of the elements predicted by Mendeleev had been isolated.

When Mendeleev's notes show that the periodic system was created in a single day, February 17, 1869. He arrived at his system by puzzling over cards containing the names of the 63 known elements along with their atomic weights and important chemical and physical properties.

Lothar Meyer-1886, also developed a periodic table based on atomic masses, independently of Mendeleev. Meyer had several inaccuracies and some elements were not included. Meyer was the first scientist to introduce the concept of valence as a periodic property. Both Mendeleev and Meyer were awarded the Royal Societies Davy Medal. Mendeleev is given credit because of his accurate property prediction of yet undiscovered elements.

Henry Moseley-1914 was a student of Rutherford. Moseley was studying X-ray formation by high energy electron bombardment. He graphed the square root of the X-ray frequency vs atomic mass. This plot gave a nearly linear line except for three atomic pairs. Ar(39.95)/K(39.10), Co(58.93)/Ni(58.69), Te(127.60)/I(126.90).

When the atoms were plotted according to atomic number, then a linear relationship was established. Moseley stated, "There is every reason to suppose that the integer that controls the X-ray spectrum is the charge on the nucleus."

Periodic Law - The properties of the chemical elements are a periodic function of atomic number.

Why Mendeleev is given Credit in Modern Text Books?

Mendeleev's Table allowed for and was capable of adjusting to future discoveries:

noble gases, new column in 1894-1901
incorporation of the rare earth elements
Moseley's atomic number in 1914
Bohr atom and electronic structure in 1913
discovery of synthetic elements 1939 to present (element 110, 1994)

The Periodic Table

Groupa vertical column of elements in the periodic table; also called a family

Perioda horizontal row of elements in the periodic table

Metalsone of a class of elements that includes a large majority of the known elements; metals are characteristically lustrous, malleable, ductile, and good conductors of heat and electricity

MetalloidsThe elements that border the stair-stepped line are classified as metalloids. The metalloids, or semimetals, have properties that are somewhat of a cross between metals and nonmetals.

Metalloids tend to be economically important because of their unique conductivity properties (they only partially conduct electricity), which make them valuable in the semiconductor and computer chip industry. The metalloids are shown in the following illustration.

Nonmetalsone of a class of elements that are not lustrous and are generally poor conductors of heat and electricity; nonmetals are grouped on the right side of the periodic table

Alkali metalsany metal in Group 1 of the periodic table. (soft, malleable, lustrous, good conductors, MOST REACTIVE family of metals)

Alkaline earth metalsany metal in Group 2 of the periodic table. (higher densities and melting points than alkali metals; not as reactive as alkali)

Halogensany member of the nonmetallic elements in Group 17 in the periodic table. ( MOST REACTIVE Non-Metals; do not occur free in nature; commonly found in sea water, minerals, & living tissues)

Noble gasesany member of a group of gaseous elements in Group 18 in the periodic table. (VERY INACTIVE elements, used in balloons, scuba diving tanks, light bulbs)

Periodic Table Exercise

The following need to be labeled on your periodic table

metals/non-metals

jewellery metals (there are three of them)

magnetic metals (three of them)

elements that are gases at room temperature

the two liquid elements at room temperature

noble gases

alkali earth elements

halogens

alkali metals

metalloids (sevenof them)

Modern Atomic Theory Notes

1850's

Robert Bunsen conducted experiments in which he observed that different elements, when heated in a flame, gave off a characteristic colour.

Late 1800's

J.J. Thomson and others were experimenting with gas discharge tubes. Gaseous elements, when subjected to electric current at low pressure, gave off a colourful glow.

1869

Dmitri Mendeleev introduced the scientific world to the idea of periodicity and that patterns of behavior within the elements were in accord with their atomic mass. Could all of these patterns follow from one property, atomic mass? Mendeleev’s periodic table began a fantastic era of scientific discovery. Scientists began an intense period of tinkering and experimentation to try and answer all the puzzling questions.

Becquerel, Curie's ---> radioactivity

Thomson, Rutherford, Chadwick, and others ---> subatomic particles

Rutherford's atomic model

Mosely ---> atomic number

Early 1900's

two extraordinary scientists, Albert Einstein and Max Planck, contributed to a significant discovery. They determined that Thomson had missed something in his study of the photoelectric effect. Thomson had shown that the negatively-charged particles emitted when a metal was struck by light were indeed the same particles that he called “electrons” from his study of cathode ray tubes.
Einstein and Planck were interested in what caused the electrons to leave an atom. They studied the phenomena of energy and light. What must happen to cause electrons to leave an atom? Energy is required to pull an electron from its attraction to the nucleus. Where does the energy come from?

Einstein - Theory of Relativity

Planck - Quantum Theory of Light

Observations:

energy in the form of heat - element gives off energy as light in a particular color.
energy in the form of electricity - element gives off energy as light in a particular color.

Light:

travels in waves with a characteristic frequency, wavelength, and energy
frequency and wavelength are inversely proportional but frequency and energy are directly proportional

In 1672, Sir Isaac Newton discovered that the diffraction of sunlight in a glass prism would produce a continuous spectrum of colors. We now know that light travels in waves so that each color travels at its own distinct wavelength.
It was later found that when one looks through a diffraction grating at an element absorbing energy and emitting light, one sees a pattern of colored lines. Each element has its own characteristic “line spectrum” which acts as a set of “fingerprints” to identify the element.

Why were elements only giving off light at certain wavelengths?

Chemistry - Wavelength, Frequency, & Energy of EMR

Show ALL equations, work, units, and significant figures in performing the following calculations.

c = λνE = hνE = hc

C = 3.00 x 108 m/sh = 6.626 x 10-34 J s

1. What is the wavelength of a wave having a frequency of 3.76 x 1014 s-1?

2. What is the frequency of a 6.9 x 10-13 m wave?

3. What is the wavelength of a 2.99 Hz wave?

4. What is the frequency of a 2,600 cm wave?

5. What is the energy of a 7.66 x 1014 Hz wave?

6. What is the frequency of a wave carrying 8.35 x 10-18 J of energy?

7. What is the frequency of a 1.31 x 10-22 J wave? What is its wavelength?

8. What is the wavelength of a 7.65 x 10-17 J wave?

9. What is the energy of a 9,330 cm wave?

10. What is the wavelength of a 1.528 x 10-13 J wave?

Chemistry – Wavelength, Frequency, & Energy of EMR

ANSWER KEY

What is the wavelength of a wave having a frequency of

3.76 x 1014 s-1?

λ = c/ν= 3.00 x 108 m/s = 3.00 x 108 m x s = 7.98 x 10-7 m

3.76 x 1014 s-1 s3.76 x 1014

What is the frequency of a 6.9 x 10-13 m wave?

ν = c/λ = 3.00 x 108 m/s = 3.00 x 108m x 1 = 4.35 x 1020 s-1

6.9 x 10-13 m s 6.9 x 10-13m

3. What is the wavelength of a 2.99 Hz wave?

λ = c/ν= 3.00 x 108m/s x 1 Hz = 3.00 x 108 m x s = 1.00 x 108 m

2.99 Hz s-1s 2.99

4. What is the frequency of a 2,600 cm wave?

ν = c/λ = 3.00 x 108 m/s = 3.00 x 108m x 1 = 1.2 x 106 s-1

2.6 x 101 m s 2.6 x 101m

5. What is the energy of a 7.66 x 1014 Hz wave?

E = h ν = 6.626 x 10-34 J/Hz x 7.66 x 1014Hz = 5.07 x 10-19 J

6. What is the frequency of a wave carrying 8.35 x 10-18 J of energy?

ν = E / h = 8.35 x 10-18J = 1.26 x 1016 s-1

6.626 x 10-34J-s

7. What is the frequency of a 1.31 x 10-22 J wave? What is its wavelength?

ν= E / h = 1.31 x 10-22J = 1.977 000 392 x 1011 s-1 = 1.98 x 1011 s-1

6.626 x 10-34J-s

λ = c/ν = 3.00 x 108m/s

1.98 x 1011 s-1

= 3.00 x 108 m x s = 0.001 52 m= 1.52 x 10-3 m

s 1.98 x 1011

8. What is the wavelength of a 7.65 x 10-17 J wave?

ν = E / h = 7.65 x 10-17J = 1.15 x 1017 s-1

6.626 x 10-34J-s

λ = c/ν= 3.00 x 108m/s =

1.15 x 1017 s-1

= 3.00 x 108 m x s =2.61 x 10-9 m

s 1.15 x 1017

9. What is the energy of a 9,330 cm wave?

ν = c/λ = 3.00 x 108 m/s = 3.00 x 108m x 1 = 3.22 x 106 s-1

9.33 x 101 m s 9.33 x 101m

E= h ν = 6.626 x 10-34 J-s x 3.22 x 106s-1 = 2.13 x 10-27 J

10. What is the wavelength of a 1.528 x 10-13 J wave?

ν = E / h = 1.528 x 10-13 J = 2.306 x 1020 s-1

6.626 x 10-34J-s

λ = c/ν = 3.00 x 108m/s = 3.00 x 108 m x 1______

2.306 x 1020 s-1 s 2.306 x 1020s-1

= 1.30 x 10-12 m

1911

Neils Bohr, a young Danish scientist working together with Ernest Rutherford, proposed a new model for the atom. The line spectrum of an element led Bohr to believe that the atom was releasing energy in the form of light only at certain “energy states.”
Bohr proposed that the electron of a hydrogen atom moves about the nucleus in a circular path of a certain radius having a certain energy state. This has been called the planetary view of the atom - electrons were found outside the nucleus in orbits moving like planets around the sun.

the lowest energy state/level is called the “ground state.” Electrons absorb energy and move from one allowed energy state to another. When electrons move to a higher energy state, they are said to in an “excited state.” When electrons fall back to lower energy states, they release energy in the form of light.

the observation that only certain wavelengths of light were absorbed or emitted led Bohr to believe that only certain energy changes were possible.
if the electron could move up to any particular energy level, than we would see a continous spectrum and not a line spectrum.
Bohr’s proposed atomic model was for the hydrogen atom (1 proton, 1 electron). His mathematical formulas and calculations for the model explained the line spectrum of the hydrogen atom. However, Bohr’s model was not able to accurately predict the line spectrum for atoms with more than 1 electron. It appeared that Bohr’s model was an oversimplification. The search to solve the mystery of the atom continued.

1923

Louis de Broglie, a French physicist, proposed that particles in motion do not travel in straight lines. Particles travel in waves!

1927

Werner Heisenberg proposed the “Uncertainty Principle.” Heisenberg reasoned that if matter, including electrons, travel in a wave-like motion then it is impossible to predict the exact path and position of an electron in the atom. Therefore, it is not correct to say that electrons move in well-defined circular orbits around the nucleus.

Late 1920's

Erwin Schrodinger, an Austrian physicist, applied mathematics to the study of an electron’s wave-like motion. This began a field of study called “wave mechanics” or “quantum mechanics.” We will not look at the mathematics involved due to its complexity, but we will look at his results and theories.
Schrodinger used probability to predict where an electron would be found in an atom at any given time. Using complex wave equations, he was able to verify Bohr’s work for the H atom and establish predictions for multi-electron atoms.
the region of space where the electron would most likely be found was called the “electron cloud.”
within the electron cloud, Schrodinger defined regions of space outside the nucleus where the electron would be found.

1.shell - main energy level

2.subshell - each main energy level is made up of 1 or more sublevels

3.orbital - each subshell is composed of 1 or more orbitals

Describing the location of the electron

Atomic orbital – three-dimensional region around the nucleus that describes the electron’s probable location

Quantum numbers

describe the location of the electron in four categories
each category gets more specific

Quantum Numbers

Energy Level (principal quantum number)

n

which can have values 1,2,3,4,5,6,7

defines the size, as n increases the energy level gets larger

Energy Sublevels (angular momentum quantum number)

each energy level has “n” number of sublevels

the sublevel have labels

1st one in each level…s

2nd…p

3rd …d

4th…f

Atomic Orbitals (magnetic quantum number)

each energy level has “n2” number of atomic orbitals

each sublevel has a fixed number of orbitals

s …. 1 orbital

p …. 3 orbitals

d …. 5 orbitals

f …. 7 orbitals

Atomic Spin (spin quantum number)

each orbital can hold a maximum of 2 electrons

each energy level has “2n2” number of electrons

Electron configurations

the arrangement of electrons in an atom
lowest energy and most stable

Rules of electron arrangement

The Rules

aufbau principle – each e – occupies the lowest energy orbital

each sublevel has a different energy state

e – within a energy level fill in sub level order…s,p,d,f

the energy levels overlap so a guideline is needed to establish sublevel order

diagonal rule – sets the order of filling the sublevels

Pauli Exclusion Principle – an atomic orbital contains a maximum of two electrons

the two e – will travel with opposite spins

direction of spin will be represented by