Chapter 6 - The Periodic Table Chapter 6 - The Periodic Table and Periodic Lawand Periodic Law

• Objectives:– Identify different key features of the periodic table.

– Explain why elements in a group have similar properties.

– Relate the group and period trends seen in the periodic table to the electron configuration of atoms

• Why this is important:– The periodic table is one of the most useful reference tools

available in chemistry. Understanding its organization and interpreting its data will aid in understanding chemistry concepts.

www.privatehand.com/flash/elements.html

Development of the Periodic Table• In 2003, there were 118 elements known.• The majority of the elements were discovered between

1735 and 1843.• How do we organize all the different elements in a

meaningful way that will allow us to make predictions about undiscovered elements?

• A variety of scientists tried to arrange the known elements to reflect the trends in chemical and physical properties but their systems did not allow for newly discovered elements to fit in their charts.

• 1869 Dmitri Mendeleev and Lothar Meyer separately arranged the elements in order of increasing atomic mass and into columns with similar properties

• That seemed to work to organize most of the elements.

http://www.chemistrydaily.com/chemistry/upload/a/a1/Dmendeleev.jpghttp://www.chemistryexplained.com/images/chfa_03_img0535.jpg

• Mendeleev is given more credit than Meyer because he published his findings first and he left spaces for elements that were not yet discovered.

• Some of the elements that he predicted were scandium, gallium, and germanium.– In 1871, Mendeleev noted that arsenic (As) properly belonged

underneath phosphorus (P) and not silicon (Si), which left a missing element underneath Si. He predicted a number of properties for this element. In 1886 Germanium (Ge) was discovered. The properties of Ge matched Mendeleev’s predictions.

• Mendeleev’s table was not completely correct.

• Arranging elements by atomic mass caused some elements to be put in the wrong groups so that the properties did not exactly match up

• 1913 English chemist Henry Moseley arranged elements in order of increasing atomic number

• Problems with order of elements were solved and there was a clear repeating pattern of properties of the elements in their groups.

• The PERIOIDC LAW states there is a “periodic” repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number.

• Periodic means: happening or reoccurring at regular intervals (definition from Webster’s Dictionary)

http://www.rsc.org/education/teachers/learnnet/periodictable/scientists/moseley.jpg

The Modern TableThe Modern Table• Boxes are arranged in order of increasing atomic #• Elements are grouped into columns by similar

properties• Scientists keep adding elements that were discovered• The final adjustment was when physicist Glenn

Seaborg had the inner-transition elements pulled below the rest of the periodic table and into 2 separate rows (This occurred in the late 1940s)

http://www.lbl.gov/Science-Articles/Research-Review/Magazine/1994/seaborgium-mag.html

• Horizontal rows are called periods• There are 7 periods

• Vertical columns are called groups or families.• Elements are placed in columns by similar

properties.– b/c of the similar numbers of valence e- they contain

1 2

13 14 15 16 17

18

3 4 5 6 7 8 9 10 11 12

The Different Groups of Elements1 – 18 system used by all chemists

A & B system is an older American System

A elements are representative elements

B elements are transition elements1A 2A

3A 4A 5A 6A 7A

8A

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

Representative or Main Group Representative or Main Group ElementsElements

• They have a wide range of physical & chemical properties.

• They have the whole range of possible valence electrons (1 to 8)

• Also called s and p block elements

• Here are some important groups:

• Group 1 (1A) contains the alkali metals (remember to NOT include hydrogen)

• Group 2 (2A) contains the alkaline earth metals

Alkali MetalsAlkali Metals

• Called this because these metals react w/ water to form alkaline (basic) solutions.

• Highly reactive metals that lose their 1 valence electron to form 1+ ions

• Soft enough to be cut with a knife.• They are stored in oil to prevent reactions with oxygen

and water in the air.

Alkaline Earth MetalsAlkaline Earth Metals

• Called this because most of these metals react with oxygen to form compounds called oxides (the alchemists called them “earths” because of this) and the oxides react w/ water to form alkaline (basic) solutions

• Not as reactive (but do react easily) & harder than group 1 metals

• They lose their 2 valence electrons to form 2+ ions

GroGroupsups 13 13 - - 1616• Named for the first element in each group.• They do have mixed groupings of elements because each

column contains nonmetals, metalloids, and metals.• Many of the elements in these groups form various

charges

• Group 17 (7A) contains the halogens• Group 18 (8A) contains the noble gases

HalogensHalogens

• Called this because halogen means “salt formers” b/c they react with metals to form salts (ionic compounds)

• Physically F & Cl are gases at room temp., Br is a liquid but it evaporates easily, and Iodine is a solid that sublimes easily

• Astatine is the odd-ball of the group b/c it’s radioactive w/ no known uses

• Chemically they are the most reactive nonmetals• They have 7 valence e- so they will share or gain 1 e-

and they tend to form 1- ions.

Noble GasesNoble Gases• Last naturally occurring elements to be discovered b/c

they are colorless & unreactive• Very stable with full valence electrons = 8

– (except He w/ 2)

• With lots of energy you can get Xe, Kr and Ar compounds (There are no known He or Ne compounds)– In 1962 the first compound of the noble gases was prepared:

XeF2, XeF4, and XeF6.

– To date the only other noble gas compounds known are KrF2 and HArF.

Transition elements (metals)

d-block

f-block

These are called the inner transition elements and they belong here

Transition MetalsTransition Metals

• Make up the majority of elements on the periodic table• Have a wide variety of uses & effect the economy• The variation of their physical properties is b/c of their

electron configurations & b/c unpaired d-electrons can move into valence shells.

• The more unpaired d-electrons, the greater the hardness & higher the melting & boiling points

• Most lose electrons to become positively charged ions• Cu, Ag, Au, Pt, and Pd are the only ones unreactive

enough to be found alone in nature

Inner Transition MetalsInner Transition Metals• Lanthanide series – follow element lanthanium

– All silvery metals w/ high melting points

• Actinide series – follow element Actinium– All are radioactive & only 3 exist in nature

HYDROGENHYDROGEN

• In a class by itself, it is a unique element• Most often occurs as a colorless diatomic gas, H2

• It is placed in Group 1 b/c it has one valence e- and will easily lose its 1 electron when reacting w/ other nonmetals to become a 1+ ion (H+ is a proton)

• But it shares many properties w/ the halogens and will sometimes gain e- when bonding w/ a metal to become a 1- ion (the hydride ion, H-)

• It is the most abundant element in the universe (90% by mass)

Metals, Nonmetals, and MetalloidsMetals, Nonmetals, and Metalloids

MetalsMetals• What are some common

properties of metals?• Have luster (shine) when smooth

& clean.

• Good conductors of heat & electricity

• Most are solid @ room temp.

• Most are:– Ductile = drawn into wires.

– Malleable = hammered into sheets.

• Most lose electrons to become cations

• Most of the elements on the periodic table are classified as metals

NonmetalsNonmetals• What are some common

properties of nonmetals?• At room temp:

– Some are brittle & dull solids

– Gases

• Poor conductors = good insulators

• Most tend to gain e- to become anions

• They have a wide variety of melting & boiling points

Metalloids or SemimetalsMetalloids or Semimetals• Chemical & Physical Properties of both metals & non-

metals• Example: Si has a metallic luster but it is brittle.• They are semiconductors b/c they do conduct electricity

but not as well as metals– Silicon (Si) & Germainum (Ge) are 2 most important for

computer chips & solar panels

Valence Electrons REVIEWValence Electrons REVIEW• These are defined as electrons in the atom’s highest energy levels.

• Examples:

• Sodium– Has 11 electrons but only one valence electron

• Cesium (Cs)– Has 55 electrons but only one valence electron

• Bromine– Has _____ electrons but only _____ valence electron

• Valence electrons help determine the chemical properties of an element and how it will bond to form compounds.

Examples – Look at the electron configuration of each element. What do you notice?

Na [Ne]3s1

K [Ar]4s1

Cs [Xe]6s1

F 1s22s22p5

Cl [Ne]3s23p5

Br [Ar]4s23d104p5

Ne 1s22s22p6

Ar [Ne] 3s23p6

Kr [Ar]4s23d104p6

Most elements in the same group have the same ending electron configuration = same number of valence electrons

• Valence electrons and groups:– Group 1 elements have 1 valence electron

– Group 2 elements have 2 valence electrons

– Group 13 elements have 3 valence electrons

– Group 14 elements have 4 valence electrons

– Group 15 elements have 5 valence electrons

– Group 16 elements have 6 valence electrons

– Group 17 elements have 7 valence electrons

– Group 18 elements have 8 valence electrons (except He, it only has 2)

– Most of the transition metals have 2 valence electrons, there are a lot of exceptions!!

Depicting Valence Electrons ReviewDepicting Valence Electrons Review• Electron dot structures – used to visually represent

valence electrons in a shorthand method.• We use an element’s symbol to show what element we

are talking about and the dots represent the atom’s valence electrons.

• In writing these structures the dots are placed one at a time on the four sides of the symbol and then paired up until all are used.

Examples of electron dot structures:

• Magnesium

• Sulfur

• Rubidium (Rb)

• Bromine

• Oxygen

Section 6.3 - Periodic TrendsSection 6.3 - Periodic Trends

Objectives:Compare period & group trends for

shielding, atomic radius, ionic radius, ionization energy, & electronegativity

Shielding Shielding (or screening)(or screening)

• What does this mean?• The valence e- are blocked from the full positive charge

of the nucleus (effective nuclear charge) by the inner (core) e-

• As the average number of core e- increases, the effective nuclear charge decreases– This idea of shielding will play a large role in a lot of the

trends

Mg

• Trend within the period (left to right): Generally decreases • Why:

– B/c the number of energy levels & core e- stays the same but the nucleus is increasing

– This increase the attraction between the nucleus and valence e-

• Trend down a group: Generally increases• Why:

– B/c the number of energy levels & core e- increases

– This makes the valence e- farther from the nucleus and more blocked by the inner e-

Shielding Shielding

Atomic RadiusAtomic Radius• ½ the distance between adjacent nuclei of identical atoms either in

crystal form (metals) or in molecular form (nonmetals)

• Trend within the period (left to right): Generally decreases

• Why: – B/c the number of energy levels & core e- stays the same but the nucleus is

increasing

– This increase the attraction between the nucleus and valence e-

– This attraction pulls the e- closer to the nucleus and makes the atom smaller

– Ex: Na vs. S

• Trend down a group: Increases• Why:

– B/c the number of energy levels increases & core e- increases

– Each energy level is larger than the next

– This makes the valence e- farther from the nucleus and more blocked by the inner e-

- EX: Na vs. K

Examples – Place each group of elements in order of increasing atomic radius:

1. S, Al, Cl, Mg, Ar, Na

2. K, Li, Cs, Na, H

3. Ca, As, F, Rb, O, K, S, Ga

Examples – Place each group of elements in order of increasing atomic radius:

1. S, Al, Cl, Mg, Ar, Na

Ar < Cl < S < Al < Mg < Na

2. K, Li, Cs, Na, H

H < Li < Na < K < Cs

3. Ca, F, As, Rb, O, K, S, Ga

F < O < S < As < Ga < Ca < K < Rb

• Ionic Radius – The distance between the nucleus and the outermost electron in ions (can’t be determined directly)

• Trend between atom & ion and Why:– Cations are smaller than original atom – b/c losing e- the atom has unequal positive charge that attracts

the valence e- closer to the nucleus

– Anions are larger than original atom and cations – b/c adding negative e- adds to the repulsion between other

valence e-, pushing them apart

Ionic Radius ContinuedIonic Radius Continued• Trend within the period (left to right): Representative

Elements– Cations the size decreases

– Anions the size drastically increases compared to the positive ions and then decreases across the period

• Trend down a group: Increases for both cations & anions

• Why: – Same reason as atomic radii trend

• For ions of the same charge, ion size increases down a group.

• All the members of an isoelectronic series have the same number of electrons.

• As nuclear charge increases in an isoelectronic series the ions become smaller:

O2- > F- > Na+ > Mg2+ > Al3+

Examples – Choose the larger species in each case:

1. Na or Na+

2. Br or Br-

3. N or N3-

4. O- or O2-

5. Mg2+ or Sr2+

6. Mg2+ or O2-

7. Fe2+ or Fe3+

• Ionization Energy: The energy required to remove an electron from a gaseous atom (also called First Ionization Energy, I1)

Na (g) + 496 kJ Na+ (g) + e-

– The second ionization energy, I2, is the energy required to remove an electron from a (1+) gaseous ion:

Na+ (g) + 4562 kJ Na2+ (g) + e-

• NOTICE: – Ionization Energy increases for each electron removed from

the same element

– The larger ionization energy, the more difficult it is to remove the electron.

Variations in Successive Ionization Energies

• There is a sharp increase in ionization energy when a core electron is removed.

• Notice the large increase after the last valence electron is removed. This chart can be used to determine the number of valence electrons in an atom of an element.

• Trends within the periods: Increases• Why: Electrons are harder to remove from smaller

atoms because they are closer to the nucleus and there is an increased nuclear charge

• Trends down a group: Decreases• Why: Electrons are easier to remove from large atoms

because they are farther away from the nucleus so there is less energy needed to remove them.

• Notice the trends in ionization energy is inversely related to trends in atomic radii.

Examples – Put each set in order of increasing first ionization energy:

1. P, Cl, Al, Na, S, Mg

2. Ca, Be, Ba, Mg, Sr

3. Ca, F, As, Rb, O, K, S, Ga

Examples – Put each set in order of increasing first ionization energy:

1. P, Cl, Al, Na, S, Mg

2. Ca, Be, Ba, Mg, Sr

3. Ca, F, As, Rb, O, K, S, Ga

1. Na < Al < Mg < S < P < Cl

2. Ba < Sr < Ca < Mg < Be

3. Rb < K < Ca < Ga < As < S < O < F

ELECTRONEGATIVITYELECTRONEGATIVITY

• Electronegativity: The relative ability of atoms that are bonded to attract electrons in the chemical bond to itself

• Chemist Linus Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). – Values are calculated from ionization energies and electron

affinities. They are used to help determine types of bonding (ionic or covalent) that are occurring in a compound.

– Noble gases are not usually given electronegativity values

• Trends within the periods: Increases

• Why: Atoms become smaller so the shared electrons are closer to the nucleus in small atoms

• Trends down the groups: Decreases• Why: Atoms become larger so shared electrons are

farther from the nucleus in large atoms

Electronegativity

Examples – put each set in order by increasing electronegativity:

1. Na, Li, Rb, K, Fr

2. Cl, Ca, F, P, Mg, S, K

Examples – put each set in order by increasing electronegativity:

1. Na, Li, Rb, K, Fr

2. Cl, Ca, F, P, Mg, S, K

1. Fr < Rb < K < Na < Li

2. K < Ca < Mg < P < S < Cl < F

Review:Review:

1. As you move across a period, left to right, describe what generally happens (decreases, increases, or remains the same) to:

a. The number of valence electrons

b. The ionization energy

c. The atomic radius

2. Give a brief explanation for your answers to a-c.

3. Identify the element from the clues given:a. This element has a smaller atomic radius than phosphorous, it

has a smaller ionization energy than fluorine and is chemically similar to iodine

b. This element has the smallest ionization energy of any element in Period 4.