Periodic Table and Electrons

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Chemistry

Learning Objectives Atomic Theory and Periodic Table

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.

CLUE ELEMENT CLUE ELEMENT 1. well drillers decision boron B 24. a spice 2. to press laundry 25. a blitz by police 3. policeman 26. …..on the Range 4. mother’s sister’s money 27. …of Arabia 5. where dishes are washed 28. dull chemistry lecture 6. a foolish prisoner 29. Golden Gate Bridge state 7. natives of North America 30. …bladder or …stones 8. water and gin 31. European country 9. shown the way 32. another European country

10. I sit down to eat 33. repair clothes 11. a Ford product 34. 50 per cent 12. a popular house plant 35. larger than a coyote 13. have the sniffles 36. God of the underworld 14. United States citizen 37. God of the sea 15. don’t take any wooden 38. a good . . . (helper) 16. playing a part or role 39. to brown a roast 17...... pop 40. 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 this helium (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

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

Period a horizontal row of elements in the periodic table

Metals one 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

Metalloids The 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.

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

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

Halogens any 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 gases any 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 (seven of 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. 13

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 = λν E = hν E = hc λ C = 3.00 x 108 m/s h = 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

1. 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 s 3.76 x 1014

2. 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 108 m x 1 = 4.35 x 1020 s-1

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

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

λ = c/ν = 3.00 x 108 m/s x 1 Hz = 3.00 x 108 m x s = 1.00 x 108 m 2.99 Hz s-1 s 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 108 m x 1 = 1.2 x 106 s-1 2.6 x 101 m s 2.6 x 101 m

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 1014 Hz = 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 10-18 J = 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 10-22 J = 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 10-17 J = 1.15 x 1017 s-1 6.626 x 10-34J-s

λ = c/ν = 3.00 x 108 m/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 108 m 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 106 s-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 1020 s-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 19

 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 1. the arrangement of electrons in an atom 2. lowest energy and most stable

Rules of electron arrangement

The Rules

1. 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

2. 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  one pair of e –

3. HUND’S RULE – e – will individually occupy equal energy orbitals before forming a pair  all orbitals of a sublevel are of EQUAL energy

Methods of notation

Writing Configurations

Orbital notation – shows every electron with an arrow

______

O 1s 2s 2p 2p 2p

Electron Configuration – shows the total number of electrons in each sublevel as a superscript.

Ar 1s2 2s2 2p6 3s2 3p6

Noble gas notation  a shorter version of electron configuration  since the inner level electron configuration doesn’t change, the noble gas is used as a shortcut

Standard version Al 1s2 2s2 2p6 3s2 3p1

same

Ne 1s2 2s2 2p6

Shorter version

Al [Ne] 3s2 3p1

Steps to use this shortcut…Cd 1. find the noble gas (group 18 on the periodic table) with an atomic number less than the element …Kr 2. Put the noble gas in brackets…Cd [Kr] 3. The noble gas filled the p sublevel with n = period number… 4 4. Follow the diagonal rule (Aufbau Principle) and continue the notation from the next s-sublevel using the remaining electrons… Cd [Kr] 5s2 4d10

ELECTRON DOT – shows each outer level electron as a dot.

Valence electrons

 electrons in the highest number energy level  found in the highest number s and p sublevels  the electrons used in electron dot notation

Atoms in the same group have similar chemical properties because they have the same number of valence electrons. Notice that the electron dot structures repeat as you move down the table.

NOTE: The electron configurations also vary by the location in the table.

Orbital diagram

Energy Level (Orbital) diagrams

These show how the electrons are placed in the orbitals as a function of increasing energy

Example: Potassium

Draw the energy level diagram for Phosphorous

Examples of writing electron configurations:

Hydrogen

Orbital notation:

Electron Configuration:

Noble Gas Configuration (shorthand notation):

Carbon

Nitrogen

Calcium

Chromium

Selenium

There are always exceptions to the above rules! Cr and Cu do not follow the rules……..

This is due to the greater stability of ½ filled or completely filled orbitals

Electron Configuration Worksheet PART A – ORBITAL NOTATION & ELECTRON CONFIGURATION Use the patterns within the periodic table to draw orbital diagrams and write longhand electron configurations for the following atoms.

- Symbol # of e Orbital notation and Electron Configuration

1. Mg

2. P

3. V

4. Ge

5. Kr

6. O

PART B –

NOBEL ELECTRON CONFIGURATION Use the patterns within the periodic table to write the shorthand electron configurations for the following elements.

- Symbol # e Shorthand (noble gas) Electron Configuration

8. Ca

9. Cu

10. F

11. Ra

PART C – RULES OF ELECTRON CONFIGURATIONS Which of the following “rules” is being violated in each electron configuration below? Explain your answer for each. Hund’s Rule, Pauli Exclusion Principle, Aufbau Principle

   __ __ 12. 1s 2s 2p

     ___  _ _ 13. 1s 2s 2p 3s 3p

        _ 14. 1s 2s 2p 3s 3p

              15. 1s 2s 2p 3s 3p 3d

ELECTRON CONFIGURATION WORKSHEET

Element symbol and FULL electron SHORTHAND electron Number of name configuration configuration- using valence electrons noble gas notation H: Hydrogen

He: Helium

Be: Beryllium

K: Potassium

Na: Sodium

P: Phosphorous

Ar: Argon

Al: Aluminum

Al3+: Aluminum

Cl-: Chloride

Cl: Chlorine

Mg2+: Magnesium

Na+: Sodium

F-: Fluoride

O2-: Oxide

Ne: Neon

Orbital Filling Worksheet

1. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element Ar.

2. Write out the electron configuration in the short-hand notation (noble gas).

3. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element Na.

4. Write out the electron configuration in the short-hand notation.

5. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element Si.

6. Write out the electron configuration in the short-hand notation.

7. Fill the boxes below with the arrow notation for electrons showing the correct ground state electron configuration for the element N.

8. Write out the electron configuration in the short-hand notation.

PERIODIC TRENDS NOTES

Periodic Law – Similar properties recur periodically when elements are arranged according to increasing atomic number.

Properties of elements – Primarily determined by the outer shell (valence) electrons.

Group 1 Metals (Alkali Metals) Outer Shell Configuration: ns1 Ion Configuration: (1+) ion is isoelectronic with a noble gas

Group 2 Metals (Alkali Earth Metals) Outer Shell Configuration: ns2 Ion Configuration: (2+) ion is isoelectronic with a noble gas

Group 13 Metals Outer Shell Configuration – ns2 np1 Ion Configuration: Al+3 is isoelectronic with a noble gas

Group 16 Non Metals Outer Shell Configuration: ns2 np4 Ion Configuration: (2-) ion isoelectronic with a noble gas

Group 17 Non Metals (Halogens) Outer Shell Configuration: ns2 np5 Ion Configuration: (1-) ion isoelectronic with a noble gas d-block elements (Transition Elements) Outer Shell Configuration: ns2/ns1 (n-1)d1-10 Ion Configuration: The electrons in the s sublevel are the first to be lost in ion formation. These ions are rarely isoelectronic with a noble gas. This is due to the presence of the electrons in the d sublevel.

Basic Assumptions

 The electrons are attracted to the nucleus o energy is required to remove an electron from an atom o energy is released when an atom gains an electron  The further away an electron is from the nucleus, the easier it is to remove o The higher the energy level, the further away the electron (think of the energy levels as shells

These are the two reasons for the  Shielding Effect (down a group) o Inner core (inner shell) electrons shield the nucleus from the outer shell electrons o Electrons in the same shell do not shield each other o This shielding is not 100% complete

 Effective Nuclear Charge (across a period) o The attractive force experienced by the outer shell electrons o Zeff ≈ (atomic number) – (# inner core electrons)

Full and half-full sublevels have greater stability

Periodic Trends Tutorial

Atomic Radius (AR)

 Group Trend: Down a group, AR increases, due to the addition of energy levels  Periodic Trend: Across a period, AR decreases, due to the increased effective nuclear charge  Shielding Effect: The result of this is that the effective nuclear charge (Zeff) experienced by the valence electrons increases across the periods. The two main assumptions of the shielding effect are: 1. The inner core electrons shield the nucleus from the valence electrons 2. The valence electrons do not shield each other note: These assumptions are not 100% accurate

Zeff ≈ (Atomic Number) – (# inner core electrons)

Ionic Radius (IR)

 Cations: Positive ions (shown in blue) are smaller than neutral atoms of the same element, the atoms have lost outer shell electrons, often the entire outer shell  Anions: Negative ions are larger than neutral atoms of the same element, the atoms have gained electrons, the nucleus cannot pull as strongly on each electron  Trends: The trends within the cations and within the anions are the same as for atomic radius

Within an isoelectronic series of ions, radii decrease with increasing atomic number b/c of increasing nuclear charge

Ionisation Energy (IE)

 The energy required to remove an electron, measured in kJ / mole  Process of removing an electron is endothermic (requires energy)  The larger the atom, the easier it is to remove an electron  General trends are the opposite of the atomic radius trends o Group Trend: Down a group, IE decreases o Periodic Trend: Across a period, IE increases  The stability of a full or ½ full sublevel causes an exception to the trend across the periods in Groups 13 and 16  Removing the second, third or fourth electron always requires more energy than removing the first electron IE1 ≤ IE2 ≤ IE3 ….

 When removing successive electrons, there is a big increase in IE when removing an electron from a full outer shell. For example, for a Group II element, IE1 ≤ IE2 ≤≤≤≤≤ IE3

Electronegativity

 The relative tendency of an atom to attract a bonding pair of electrons when the atom is chemically combined with another atom  High EN (non metals) – tendency to attract the bonding pair (often win the “tug of war”)  Low EN (metals) – lower tendency to attract the bonding pair (often loses the “tug of war”)  General trends are the opposite of the atomic radius trends

o Group Trend: Down a group, EN decreases

o Periodic Trend: Across a period, EN increases

Physical and Chemical Properties

Non-metals:  Melting point, boiling point, and the state of matter at room temperature change gradually within a group o Group Trend: Down a group, melting point and boiling point increases, state goes from gas → (liquid) → solid o Periodic Trend: No predictable trend  Reactivity decreases down a group  Metallic character increases down a group Metals:

 Metallic Character o Group Trend: Down a group, metallic character increases o Periodic Trend: Across a group, metallic character decreases  Reactivity increases down a group

PERIODIC TRENDS WORKSHEET

1. Choose which statement about the alkali metals lithium and cesium is correct. a) as the atomic number increases, the Electronegativity of the elements increases b) as the atomic number increases, the melting point of the elements increases c) as the atomic number increases, the first ionization energy of the elements decreases d) as the atomic number increases, the atomic radius decreases e) as the atomic number increases, the electron affinity increases

2. The following elements and ions are isoelectronic. Determine which of the following shows the correct order of their increasing radii: a) K+> Ar>Ca2+ b) Ar>K+> Ca2+ c) Ca2+>K+>Ar d) Ca2+>Ar>K e) They all have the same radii

3. Determine which element you would expect to have the lowest first ionization energy. Li Cs H He Ba

4. Identify which atom should have the largest value for the electron affinity: He F Na Si Mn

5. The following is a list of the usual charge found on the ions of a series of elements: A- G2+ Z2- V3+ X+

State which elements are most likely to be metals:

a) V,G and X b)V and G c) X and A d) A and Z e) only X

6. Describe the relationship between the group number and the electron configuration of the elements in a group.

7. Arrange the following elements in order of decreasing atomic size: sulfur, chlorine, aluminum and sodium. Explain if your arrangement demonstrates a periodic trend or a group trend

8. Indicate whether the following properties increase or decrease from left to right across the periodic table. Account for the trend using he atomic model. a) atomic radius (excluding noble gases) b) first ionization energy c) electronegativity d) metallic character

9. Describe the relationship between a) ionisation energy and atomic radius b) ionisation energy and electronegativity

10. Would you expect the ionization energies for two isotopes of the same element to be the same or different? Justify your answer.

11. When a chlorine atom forms an ion its radius increases, but when a sodium atom forms an ion its radius decreases. Explain this apparent contradiction.

Worksheet 2

Use a periodic table to help you answer the following questions. 1. Which element in the second period has the greatest atomic radius? 2. Which of the group IIIA (13) elements is the largest? 3. Of the halogens, which has the smallest radius? 4. Which of the alkaline earth metals is the largest? 5. Which of the transition metals has the smallest atomic radius? 6. Which of the noble gases is the smallest? 7. The atomic radius of which element is the largest? 8. Do alkali metals generally make anions or cations? 9. Which of the elements which have their valence electrons in the second energy level is the largest? 10. Which of the metalloids has the smallest atomic radius? 11. Which of the rare earth elements is the smallest? 12. Which of the transition metals in the fifth period is the largest? 13. Are metal ions larger or smaller than the neutral atoms they came from? 14. Are cations larger or smaller than the neutral atoms they came from? 15. Are ions of alkali metals larger or smaller than ions of alkaline earth metals from the same period? 16. Which element in the second period has the greatest first ionization energy? 17. Which of the group IIIA (13) elements has the largest ionization energy? 18. Of the halogens, which has the smallest electronegativity? 19. Which of the alkaline earth metals has the smallest electronegativity? 20. Which of the transition metals has the largest ionization energy? 21. Which of the noble gases has the smallest ionization energy? 22. Which of the group IVB (14) metals is the least active? 23. Which of the halogens is the most active? 24. Which of the semi-metals that have their valence electrons in the fourth energy level has the largest ionization energy?

25. Which of the period three elements has the largest electronegativity? 26. Which of the inner transition elements of the seventh period is the easiest to ionize? 27. Which of the transition metals in the fifth period has the largest EN? 28. Which of the group four metals has the largest ionization energy? 29. Which of the non-metals in the third period is the most active? 30. As atomic size increases, what happens to the ionization energy of the atom?

Worksheet 3

Use the periodic table and your knowledge of periodic trends to answer the following questions.

Which atom in each pair has the larger atomic radius? a) Li or K b) Ca or Ni c) Ga or B d) O or C e) Cl or Br f) Be or Ba g) Si or S h) Fe or Au

What is the periodic trend for atomic size from top to bottom in a group? from left to right in a period?

Why do atoms get smaller as you move left to right in a period?

Which element in each pair has a larger ionization energy? a) Na or O b) Be or Ba c) Ar or F d) Cu or Ra e) I or Ne f) K or V g) Ca or Fr h) W or Se

Explain the relationship between the relative size of an ion to its neutral atom and the charge on the ions.

Which particle has the larger radius in each atom/ion pair? a) Na, Na+ b) S, S2- c) I, I- d) Al, Al3+

What is ionization energy? What is first ionization energy?

What is the periodic trend for first ionization energy?

Arrange the following groups of elements in order of increasing ionization energy. a) Be, Mg, Sr b) Bi, Cs, Ba c) Na, Al, S

Which element in each pair has a higher electronegativity value? a) Cl, F b) C, N c) Mg, Ne d) As, Ca

What is the periodic trend for electronegativity?

Worksheet 4

1. Rank the following in order of increasing ionic radius: O, S, and F.

2. Which has the largest ionization energy: N, O, or Cl?

3. In each of the following pairs, circle the species with the largest electronegativity.

a. Li or Cs b. Cl or Ar c. Ca or Br d. Na or Ne e. B or Be

4. Circle the best choice in each list:

a. highest ionization energy: C, N, Si

b. largest ionic radius: P-3, S-2, Cl-1

c. highest electronegativity: As, Sn, S

d. smallest atomic radius Na, Li, Be

e. least reactive Rb, K, Be

5. In each of the following pairs, circle the species with the larger atomic radius:

a. Mg or Ba b. S or S-2 c. Cu+2 or Cu d. He or H e. Na or Cl

6. In each of the following pairs, circle the species that is most reactive.

a. Li or Be b. Fe or Cu c. P or Cl d. Na or Fr e. O or Cl

7. Indicate whether the following properties increase or decrease from left to right across the periodic table. Account for the trend using the atomic model.

Trend Increasing or Decreasing Explanation Atomic radius

Ionization energy

Electronegativity

Worksheet - An Alien Periodic Table

Background Information: Earth's scientists have announced that they have made radio contact with intelligent life on a distant planet. One of this alien planet's languages has been translated and scientific information has begun to be exchanged. The planet is composed of the same elements as Earth. However, the inhabitants of the planet have different names and symbols for them. Since the alien scientists do not know the names of our elements, they have radioed the following data on the known properties of the elements. Their periodic table only consists of the main group of elements: 1, 2, 13, 14, 15, 16, 17, & 18. The data are as follows:

Procedure: Fill in the blank periodic table with the correct alien planet symbol for each element. The symbol is given in parentheses after the element name in the data statements.

8. The inert gases are Bombal (Bo), Wobble (Wo), Jeptum (J) and Logon (L). Bombal (Bo) is a noble gas but does not have 8 valence electrons. The outside energy level of Logon (L) is its second energy level. Of these noble gases, Wobble (Wo) had the greatest atomic mass. 9. The alkali metals are Xtalt (X), Byyou (By), Chow (Ch) and Quackzil (Q). Of these alkali metals, Chow (Ch) has the lowest atomic mass. Quackzil (Q) is in the same period as Wobble (Wo). 10. The halogens are Apstrom (A), Vulcania (V), and Kratt (Kt). Vulcania (V) is in the same period as Quackzil(Q) and Wobble (Wo). 11. The metalloids are Ernst (E), Highho (Hi), Terriblum (T), and Sississ (Ss). Sississ (Ss) is the metalloid with the highest atomic mass. Ernst (E) is the metalloid with the smallest atomic radius. Highho (Hi) and Terriblum (T) are in group 14. T has more protons than Hi. The element Yazzer (Yz) is a metalloid by location but has properties that suggest it is a light metal. 12. The metallic element with the greatest atomic radius is called Xtalt (X). 13. The non-metal with the highest electronegativity on the planet is called Apstrom (A). 14. The lightest element on the planet is called Pfsst (Pf). 15. The element that forms a 4 ion with the lowest ionization energy is Elrado (El). 16. The chemical makeup of the alien planet’s oceans seems to be about the same as Earth's oceans. When seawater is distilled, the liquid that is boiled off and then condensed has been shown to have molecules consisting of two atoms of Pffst (Pf) and one atom of Nuutye (Nu). The solid left behind after distillation consists mainly of a crystal made up of the elements Byyou (By) and Kratt (Kt). 17. The element called Doggone (D) has only 4 protons in its atom. 18. Floxxit (Fx) is a black crystal and has 4 electrons in its outer most energy level. 19. Both Rhaatrap (R) and Doadeer (Do) have atoms with four energy levels. But Rhaatrap is less metallic than Doadeer. 20. Magnificon (M), Goldy (G), and Sississ (Ss) are all members of group 15. Goldy has fewer total electrons than Magnificon. 21. Urrp (Up), Oz (Oz), and Nuutye (Nu) all gain 2 electrons. Nuutye is diatomic. Oz has a smaller radius than Urrp. 22. The element Anatom (An) tends to lose 3 electrons. 23. The elements Zapper (Z) and Pie (Pi) both lose 2 electrons. Pie has lower electronegativity than Zapper.

1 The Alien Periodic Table 18

2 13 14 15 16 17

Periodic Table Puzzles

Fictitious symbols are used for the first 18 elements in the periodic table. Use the clues below to write the fictitious symbol in the appropriate spot on the periodic table provided. Symbols for real elements do not represent those elements. HINT: You do not have to complete each clue in order.

Puzzle #1:

Clue 1 U and J are alkali metals. J has more energy levels. Clue 2 T has 4 valence electrons on the 3rd energy level. Clue 3 M is a metal in period 3 with 2 valence electrons. Clue 4 X has one proton in its nucleus. Clue 5 Q has 2 energy levels, is a nonmetal, and is a solid at room temperature. Clue 6 L is a noble gas that doesn’t have 8 valence electrons. Clue 7 Z and Y are members of the nitrogen family. Y is a gas at room temperature. Clue 8 D has an ending electron distribution of s2p5. R has an ending electron distribution of s2. Clue 9 G has 6 valence electrons. Clue 10 V and W have full outer energy levels. V has 3 energy levels. Clue 11 A atoms have 3 valence electrons and E atoms have 6 valence electrons. Both are in the second period. Clue 12 K has one fewer total electrons than V. Clue 13 I has 3 valence electrons on the third energy level.

Puzzle #2:

Clue 1 Lg has 5 valence electrons on the second period. Clue 2 Eg atoms have 12 protons in the nucleus. Clue 3 Qp and Ju are halogens. Ju has fewer energy levels. Clue 4 Ke is a member of the oxygen family. Ke is in the same period as Lg. Clue 5 Gn is a member of the nitrogen family. Clue 6 Rm and Sk have 3 valence electrons. Rm has more occupied energy levels than Sk. Clue 7 Td and Vo are metals in the same family. Vo has 2 energy levels and Td has 3 energy levels. Clue 8 Wa is a member of the alkaline earth metals. Clue 9 Zy has an ending electron distribution of 3p4. An has 18 total electrons. Clue 10 Ms is a nonmetal located on the side of the periodic table with all of the metals. Clue 11 Bx atoms and Oz atoms are stable. Oz atoms are heavier than Bx atoms. Clue 12 Ds and Cy are members of the carbon family. Ds has fewer protons than Cy.