Chemistry Test: Periodic Table

The early periodic table: 1700s

Only 30 identified elements up to this time

The early periodic table: 1800s

Scientist discovered another 30 elements through atomic spectroscopy (light analysis like with gas tubes)

The early periodic table: Johann Döbereiner

• 1829

• German pharmacy apprentice

• He classified elements into triads based on mass averages

The early periodic table: John Newlands

• 1863

• English analytical chemist

• He arranged the elements into octaves (law of octaves) by increasing atomic mass

  • Colleagues didn’t take it seriously because of relationship between music and chemistry

Dimitri Mendeleev

• He was looking for patterns he could use to organize the elements in a logical way

• He observed that chemical and physical properties of elements followed a periodic pattern

What did Mendeleev observe that led to the periodic table

• Mendeleev observed repeating patterns in physical properties such as density

• When density versus atomic mass in graphed, Mendeleev found that the density of the elements rose and fell in a repeating cycle

What is a periodic pattern

At regular intervals, a pattern would repeat

The modern periodic table: Henry Moseley

• 1913

• English physicist

• He developed Moseley’s law which showed a relationship between atomic wavelength and atomic number

• This caused the “shuffling” of the periodic table to be organized by atomic number and not atomic mass

The modern periodic table: Periodic Law

When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern

The modern periodic table: metal-nonmetal line

The dark “staircase” line on the periodic table divides the table into two sections of metals and non-metals

What are all the families on the periodic table

• Alkali metals

• Alkaline earth metals

• Transition metals

• Rare earth metals

• Non-metals

• Halogens

• Noble gasses

Information on the periodic table

• Each horizontal row on the periodic table is called a period

• Each column is called a group

• The number beneath each element symbol is the atomic mass

Metalloids/Semi-metals

• Located along the metal-nonmetal line

• Properties: Combinations of both metals and nonmetals

What elements should be heavily outlines as metalloids

Boron, silicon, germanium, arsenic, antimony, tellurium, and astatine

Metals

• Located to the left of the staircase line

• Metallic properties increase as distance from staircase increases

What are the properties of metals

• Have luster

• Good conductors

• Malleable

• Ductile

• Typically solid at room temperature (except Hg)

Nonmetals

• Located to the right of the staircase line

• Don’t have metallic properties

• Typically gases at room temp (some solids - Br is a liquid)

• Variety of physical properties (many colors; soft/hard solids)

Organization of the periodic table

• Elements that share physical and chemical characteristics are place together in groups

• Each is called a family because family members have similar physical and behavioral traits

What do electron configurations describe

How electrosn are distributed among s, p, d, and f orbitals

The noble gases

• Noble gases have full outermost s and p orbitals

• These outermost electrons determine chemical behavior

• Electrons in completely filled energy levels are not available to form chemical bonds

• Electrons in unfilled energy levels are available to form chemical bonds

Valence electrons

• Electrons that are in the highest energy levels are called valence electrons

• Valence electrons participate in chemical bonds

• Valence electrons are lost or gained in ionic bonds and are shared in covalent bonds

The octet rule

• The rule that compounds formed so as to give every atom 8 valence electrons is called the octet rule

• 8 valence electrons completely fill the s and p orbitals

• Full energy levels are very stable

Light elements

• The octet rule is really a “duet” rule for hydrogen, helium, lithium, and beryllium

• These elements are isoelectronic with He and only need 2 electrons to complete their energy levels

Alkali metals

• The alkali metals (group 1) all lose an electron easily

• They form +1 ions.

• Once they lose an electron they have full energy levels.

Properties of alkali metals

• Malleable

• Ductile

• Good conductors

• Very reactive

Alkaline earth metals

• The alkaline earth metals (group 2) all lose an electron easily

• They form +2 ions

• Once they lose two electrons they have full energy levels

Properties of alkaline earth metals

Reactive but not as much as alkali metals

The halogens

• The halogens (group 17) all gain an electron easily

• They form -1 ions

• Once they gain an electron they have full energy levels

Propertes of halogens

Very reactive, especially with alkali metals and alkaline earth metals

Carbon, nitrogen, and oxygen

• Carbon, nitrogen, and oxygen are among the most important elements to life

• Oxygen and nitrogen both tend to accept electrons, rather than donate them

• Carbon can go either way, sometimes accepting electrons and sometimes donating them.

Transition metals

• The transition metals are the elements in the middle of the periodic table

• The transition metals all have electrons in d orbitals

Lanthanides

• Silvery white metals that tarnish when exposed to air

• They have high melting and boiling points

• Very reactive

Actindes

• All reactive

• Combine directly with most non-metals

• Very dense metals with distinct structures

Atom Size, Valence, & Reactivity

• As you move from period 2 to period 7 down the group 1 alkali metals, reactivity increases because:

  • Alkali metals only need to release one valence electron in order to achieve a full energy level, AND

  • Atomic size also increases as you go down a group; THUS

  • A single valence electron is more easily removed as atomic size increases

General group trend

• As you move down a group, additional energy levels are added

• They take up space that makes electrons move to significantly farther away positions from the nucleus

• As electron and proton attractions vary over distance, this “loosens” the electrons at the outer reaches of the atom, thus making them less tightly held, and less attracted to the nucleus

Shielding effect

• The additional layers of electrons that are added as you move down a group “block” the next level from being more attracted to the nucleus too

• The shielding effect increases down any group (so as you move down, the outer electrons are less tightly held)

General period trend

• As you move from one element to the next across a period on the periodic table, one additional proton and one additional electron are added

• The nucleus become increasingly “powerful” and the electrons are held tighter to the nucleus

Effective nuclear charge

• Because of the repulsion among electrons, they “lose” to the protons meaning that the addition of each proton strengthens the nucleus much more than the addition of an electron

• The nucleus becomes increasingly powerful with each additional proton, and the electrons become held even tighter around the nucleus

• The strengthening of the nucleus across the period is called the effective nuclear charge

What is the atomic radius

• The distances measured from the center of the nucleus to the outer reaches of the electron cloud

• This is impossible to measure so instead we take the distance between two nuclei that are combined and divide by two

  • This works as long as atoms are measured in the same way

Shielding effect group trend

Shielding effect increases down a group → valence electrons less tight/more spread out → increases radius

Effective nuclear charge period trend

Effective nuclear charge increases across a period → hold electrons tighter to nucleus → decreases radius

Ionization energy

• The energy required to remove the “last” electron from an atom

• The closer and more tightly held an electron is, the higher this energy will be

• There are some blips in the ionization energy trends

Ionization energy group trend

Down a group, the ionization energy will decrease as valence electrons are less tightly held (shielding effect)

Ionization energy period trend

Across a period, the ionization energy increases as the radius contracts (effective nuclear charge)

Cation

An ion with a positive charge

Anion

An ion with a negative charge

Ionic radius periodic trends: Cations

• Cations are smaller than the atoms from which they came - the loss of electrons means that the same number of protons are holding fewer electrons

• Therefore, the protons can hold the electrons more efficiently and the radius decreases

Ionic radius periodic trends: Anions

• And ions are larger than the atoms from which they came - the gain of electrons means that the same number of protons are holding more electrons

• Therefore, the protons can hold the electrons less efficiently, and the radius increases

Ionic radius trend

• The ionic radius trend is not as neat and tidy as the other trends, there appears to be a break

• This occurs where the metals (cation formers) and the nonmetals (anion formers) start

Coulomb’s law

• Matter can be electrically charged and the charges can be negative, positive, or neutral

• Like (same) charges repel and opposite charges attract

• The closer two charges are the stronger the force between them

• Coulomb’s law relates the force between two object objects to their electrical charges