PERIODIC TABLE DEVELOPMENT



SC4. Students will use the organization of the Periodic Table to predict properties of elements.

a. Compare and contrast trends in the chemical and physical properties of elements and their placement on the Periodic Table.

PERIODIC TABLE

DEVELOPMENT

DMITRI MENDELEEV mid 1800’s

- discovered a basic chemistry principle

- felt there was a certain pattern with the elements

- tested his hypothesis that there was a periodic relationship among the elements.

- set up the periodic table by atomic mass.

- left blanks for undiscovered elements (3 were later discovered).

HENRY MOSELEY 1913

- some elements were out of place in Mendeleev’s table.

- determined atomic numbers using x-rays.

- elements were placed according to atomic number. This was an important change.

- This is the modern periodic table.

PERIODIC LAW

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

THE PERIODIC TABLE

Columns vertical, called groups or families

18 total (8 main ones)

elements are not identical, but similar in properties

Rows horizontal, called periods, 7 total

elements are not alike

PATTERN: left side elements are active solids, far right side elements are inert gases. Last two rows are rare earth elements. Atomic # increases from left to right.

Valence Electrons (The electrons available for bonding)

Pattern:

Column 1 (IA) has one

Column 2 (IIA) has two

Column 13 (IIIA) has three

Column 14 (IVA) has four

Column 15 (VA) has five

Column 16 (VIA) has six

Column 17 (VIIA) has seven

Column 18 (VIIIA) has eight

Electron Configuration: Columns 1- 2 (IA - IIA) are the ‘s’ orbital.

Columns 13-18 (IIIA-VIIIA) are the ‘p’ orbitals (not He)

Columns 3-12 (IB-VIIIB) are the ‘d’ orbitals.

Lantanides and Actinoids are the ‘f’ orbitals.

METALS

• left side

• good conductors of heat and electricity

• hard and shiny (not always silver)

• can be pounded into different shapes - malleable

• can be drawn into a wire - ductile

• high density, high melting points

• react with water and substances in the atmosphere

(ex. rusting, tarnishing)

• has only a few electrons in the outer level. General rule: 3 or fewer electrons in outer level are considered to be metals. Metals have a tendency to lose these electrons when forming compounds

• most elements are metallic

NON-METALS

• right side

• poor conductors of heat and electricity (solids are insulators)

• brittle solids or gases

• dull, shatter easily

• lower density, lower melting points

• not as easy to recognize as a group

• has more than 4 electrons in the outer level. General rule: 5 or more electrons in outer level are considered to be nonmetals. Nonmetals have a tendency to gain electrons when forming compounds

METALLOIDS

• have properties of metals and nonmetals

• “metal-like”

• located on either side of the staircase

• all are shiny, white-gray in color

• all are solids

• okay conductors, ductile, malleable

SC4. Students will use the organization of the Periodic Table to predict properties of elements.

b. Use the Periodic Table to predict periodic trends including atomic radii, ionic radii, ionization energy, and electronegativity of various elements.

PERIODIC TRENDS

From Left to Right: nucleus becomes more positive, more attraction occurs, elements become more non-metallic

Within a group: the elements become more positive, but there is a larger nucleus due to the increased number of protons and neutrons. Therefore, there is less attraction between protons and outer electron levels and have more metallic properties.

Remember that elements are arranged according to atomic number. Both the position and the properties of the elements arise from their electron configuration (which comes from their atomic number).

Same column = similar outer level e- configuration

Properties that are periodic:

metals, metalloids, nonmetals, boiling point, density, atomic radii, ionic radii, oxidation numbers, ionization energies, electron affinities, electronegativity

VIDEO: Periodic Trends: Reactivity

What is the chemical equation for the reaction for each of these metals in water?

Li + H2O (

Na + H2O (

K + H2O (

SHIELDING

The higher the principle quantum number, the more energy levels there are. The orbitals are further out. Inner electrons shield the positive charge of the protons and the electrons in the outer energy level come off more easily.

PROTON PULL (also called nuclear charge)

The greater the positive charge (due to increased atomic number), the greater the attraction within the energy level.

ATOMIC RADII page 188

As the energy levels increase, the principle quantum number increases , so does the atomic radii.

As the atomic number increases across a period, the positive charge also increases within the same energy level.

TREND: atomic size increases down a group, atomic size decreases across a period.

Why: within a group, inner level electrons shield outer level e- from the positive nucleus; distance from nucleus increases; within a period, size of nuclear charge increases and the attraction between protons and neutrons pulls the atom in tighter.

IONIC RADII page 190

When atoms form a compound, the compound is more stable than the uncombined atoms were. The outer energy levels are full (become like the noble gases).

TREND: Metallic ions are smaller than the atoms they come from. Nonmetallic ions are larger than the atoms they come from.

Why: metal lose electrons and become the noble gas configuration on the energy level below. Nonmetals gain electrons and become the noble gas configuration at the end of their period.

Example: Sodium, Na Chlorine, Cl

OXIDATION NUMBERS

This is the charge number an atom would have if the valence electrons were completely transferred/lost.

Group 1 (1A) 1+ oxidation (Hydrogen can have 1+ or 1-)

Group 2 (IIA) 2+ oxidation

Group 3-12 (IIIB) can have 1+ to 8+ (lose “d” and then “s”, one at a time)

Group 13 (IIIA) 3+

Group 14 (IVA) 2+ or 4+

Group 15-17 gain electrons (-3, -2, -1)

Group 18 noble gases; 0

IONIZATION ENERGY page 191-192

This is the energy required to remove an electron. The first ionization energy is the energy required to remove the most loosely held electron.

TREND: Ionization energy increases as atomic number increases across a period. In a group, there is a gradual decrease in the first ionization energy as the atomic number increases. Metals have low first ionization energies and nonmetals have high first ionization energies.

Why? Across a period, the larger the nuclear charge, the greater the ionization energy; Down a group, the greater shielding effect leads to less ionization energy; With a bigger electron cloud (more energy levels), the ionization energy decreases; an electron in a full or half-full sublevel requires additional energy to be removed.

ELECTRON AFFINITY

The attraction of an atom for an electron; the energy change that occurs when an atom gains an extra electron. The same factors that affect ionization energies affect electron affinities.

TREND: Nonmetals have large electron affinities (except for Noble gases). The more stable an atom is, the less the tendency to have an affinity for electrons. Metals have low electron affinities and Nonmetals have high electron affinities. Nobel gases (p6) and Alkaline Earth Metals (s2) have negative affinities (extremely low).

ELECTRONEGATIVITY page 194

A tug of war between atoms for electrons in a chemical bond. How well the electrons “tug” is electronegativity. The better the atom ‘tugs’ an electron, the higher the electronegativity

TREND: increases from left to right across a period as the number of valence electrons increases and the size of the atom decreases. A low ionization energy and a low electron affinity means low electronegativity. Fluorine is the most electronegative element.

STABILITY

No

special

arrangement

half-full sublevel

full sublevel

full outer energy level

SC4. Students will use the organization of the Periodic Table to predict properties of elements.

c. Compare and contrast trends in the chemical and physical properties of elements and their placement on the Periodic Table.

THE ELEMENT FAMILIES

Special Case

HYDROGEN - acts as a metal and a nonmetal; can lose or gain its electron; most abundant element in the universe, flammable

Group One (IA) ALKALI METAL FAMILY

- very reactive, lose their electron more easily as you go down the family. Most reactive is francium

- not found in nature by themselves, form 1+ ions, 1 valence electron

- low ionization energy, low electron affinity, low electronegativity

- SODIUM, POTASSIUM important to body functions

- LITHIUM is exceptional due to its small size; more like magnesium due to diagonal trend

- FRANCIUM – most reactive metal, but extremely rare

Group Two (IIA) ALKALINE EARTH METAL FAMILY

- less reactive, but similar to alkali metals

- forms 2+ ions, 2 valence electrons

- low ionization energy, low electron affinity, low electronegativity

- CALCIUM – 5th most abundant element on earth (Lime, calcium chloride, body functions)

Groups Three - Twelve TRANSITION ELEMENTS

- Includes elements that do not fit into other groups; “d” block

- They all have properties similar to one another and to other metals: resistant to corrosion, high melting points, brittle

- These metals are less reactive than Group 1 or 2, and are harder.

- They chemically combine with oxygen to form oxides.

- They usually have 2, 3, or 4 valence electrons and form positive ions

- Include IRON (steel), CADMIUM (batteries), COPPER (wiring), COBALT (magnets), SILVER (dental fillings), ZINC (paints), GOLD (jewelry)

Group Thirteen (IIIA) BORON FAMILY

- BORON – metalloid; more like silicon, then its own family

- ALUMINUM – most plentiful metal in the earth's crust, has the most practical uses

- GALLIUM – low melting point, component of blue lasers.

Group Fourteen (IVA) CARBON FAMILY

- generally react by sharing electrons, is a 4+ or 4- ion

- CARBON – most versatile element can form millions of compounds; field of organic chemistry; has several allotropes: graphite, diamond, fullerine, carbon black

- SILICON – second most plentiful element in the earth’s crust (quartz); many industrial uses.

- LEAD – toxic; used to be in paint, plumbing, gasoline

Group Fifteen (VA) NITROGEN FAMILY

- five electrons to share, forms a 3- ion

- NITROGEN – found in fertilizers, TNT, medicines, proteins

- PHOSPHORUS – compounds found in laxatives, cheese, and baking powders

Group Sixteen (VIA) OXYGEN FAMILY

- gains two electrons, 2- ion

- OXYGEN - very reactive; most plentiful element in the earth’s crust; forms compounds with practically every element (except neon, argon, and helium); has two allotropes, O2 and O3.

Group Seventeen (VIIA) HALOGEN (salt forming) FAMILY

- Reacts with Alkali Metals to form salts (KCl)

- gain one electron to become a negatively charged ion, 1-, 7 valence electrons

- high ionization energy, high electron affinity, high electronegativity

- FLUORINE - most reactive nonmetal (element) due to its size; most electronegative element; reacts with all elements but neon, helium, and argon.

- CHLORINE – deadly gas; compounds act as bleaching agents and disinfectants

- IODINE – used to disinfect water and wounds (tincture of iodine)

Group Eighteen (VIIIA) NOBLE GAS FAMILY

- Non-reactive, have stable outer electron configurations

- All electron energy levels are full, have 8 valence electrons

- high ionization energy, low electron affinity, low electronegativity

- HELIUM – used in scuba diving; balloons

- NEON, ARGON – lighting

INNER TRANSITION ELEMENTS

all elements in the “f” block.

Lanthanoids:

- silvery metals with high melting points

- high luster and conductivity

- hard to separate since they are so similar

- used in making high quality glass, television screens, lasers, tinted sunglsses

Actinoids:

- all are radioactive

- 93-103 are synthetic (called transuranium elements); are created in particle accelerators

- Most Common are URANIUM and PLUTONIUM - used as nuclear fuel.

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LEAST ENERGY

GREATEST

ENERGY

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