Chemistry Quest #2

Periodic Table Trends, Atomic Mass, Isotopes

How to calculate atomic mass

the # of protons + # of neutrons = atomic mass amu

Why decimal places?

• Atomic mass of a single atom is a whole number; Element masses are based on Carbon-12 = 12 amu; Hydrogen (1 amu) is 1/12 the mass of carbon

What is an isotope

An isotope is a version of an element with a unique mass and number of neutrons

Isotopes

Nearly every element on the periodic table has at least two isotopes

Relative abundance

Isotopes have varying relative abundance: they are found in nature in differing amounts

Trace

the isotope is found in very small amounts

Calculating weighted average

Steps of calculating weighted average

1. Set equation with subscripts

2. Sub in values

3. Final answer and units

Radioisotopes

A radioactive isotope

Properties of radioisotopes

• Has a very unstable nucleus

• Can occur naturally but many are created artificially

• When the nucleus of a radioisotope decays, energy is released in the form of radiation

Atomic radius

the distance from the centre of the atom (nucleus) to the valence shell

Nuclear charge

• The number of protons in the nucleus determines its positive charge.

• If the number of protons changes, the strength of the nucleus’s pull on the electrons also changes.

Electron shells

• Electrons only exist where electron shells or energy levels allow them to

• A difference in the number of energy levels affects the distance electrons can be from the nucleus

Shielding

• Electrons are attracted to protons

• Inner electron shells bloc, or shield the attractive force between outer electrons and the nucleus

Electron-electron repulsion

• Electrons in the same electron shell repel each other (same charge)

• although, there is a greater effect due to the increased number of protons in the nucleus

TIPT - going across a period

• Increased number of protons in the nucleus

• Number of shells and shielding between the nucleus and valence electrons remain constant

• Atomic radius decreases

TIPT - going down a group

Despite the increasing number of protons, the atomic radius increases down a group because added electron shells and greater shielding reduce the nucleus’s pull on valence electrons.

What is a trend in the periodic table

A trend shows a general pattern, but there can be exceptions, making it different than a rul

Ionic raidus

• Ideally, the distance from the centre of an ion to its valence shell

• It is difficult to measure consistently

Ionic radius in metals

Metals tend to lose all of their valence electrons to form cations

Ionic radius in non metals

• non metals tend to gain electrons to fill their valence shell to form anions

  • same number of protons and electron shells, same shielding

  • different number of electrons resulting in more electron repulsion, and small force

Cation and anion size

• A cation is smaller than its parent atom because it has fewer electron shells and less shielding, so the nucleus pulls the electrons closer.

• An anion is larger than its parent atom

• Cations are smaller than most anions

Ionic radius - trends across a period

• Cations: Ionic radius decreases across a period.

  • Valence electrons are in the same energy level.

  • Shielding stays the same, but protons increase → stronger attraction, smaller ions.

  • Cations are often isoelectronic (same electron arrangement).

• Anions: Ionic radius also decreases across a period.

  • Valence electrons in the same shell, shielding unchanged.

  • More protons pull electrons closer.

  • Anions are also isoelectronic with each other.

electron affinity 

• Ideally, a measure of the attractive force of an atom has for adding an electron to its valence shell

Nuclear charge - electron affinity

More protons in the nucleus = greater attraction to an additional electron

Electron shell/shielding (electron affinity)

• Greater distance = decreased attraction

• Greater shielding = decreased attraction

• between nucleus and incoming electron

Electron affinity across a period

• Increasing number of protons in the nucleus increases attraction between the nucleus and incoming electrons

• number of shells and shielding between the nucleus and incoming electrons remain constant

• An increasing amount of shielding between the nucleus and the incoming electron reduces the attractive force

First ionization energy

• A measure of the energy required to remove ONE electron from the valence shell of an neutral atom

Nuclear charge - first ionization energy

• More protons in the nucleus = stronger attraction to valence electrons.

• Stronger attraction = more energy required to remove an electron.

electron shells/shielding - first ionization energy

• More electron shells = valence electrons farther from nucleus.

• Greater distance + more shielding = weaker attraction.

• Weaker attraction = less energy needed to remove an electron.

FIE - Trends going across a period

• Increases across a period.

• More protons → stronger attraction to electrons.

• Same shells & shielding → no added distance or blocking.

• Result: harder to remove an electron.

FIE - down a group

• Decreases down a group.

• More shells + shielding → valence electrons farther from nucleus.

• Weaker attraction → easier to remove electrons.

• Despite more protons, shielding dominates.

Electronegativity

• Electronegativity = an atom’s ability to attract shared electrons in a covalent bond.

• It is a relative measurement (compared to other elements).

• Determines bond type:

  • Pure covalent (equal sharing)

  • Polar covalent (unequal sharing)

  • Ionic (electron transfer)

• Greater nuclear attraction to shared electrons = greater electronegativity.

Factors affecting electronegativity

• nuclear charge 

• electron shells/shielding

Electronegativity - going across a period

• increasing electronegativity

• increase of protons in the nucleus

• greater hold and attraction on the shared valence electrons

• electron shielding/shells stay the same

E

Electronegativity - down a group

despite the increase of protons in the nucleus, electronegativity decreases since there is an increase of electron shells and shielding between the nucleus and shared valence electrons, causing less of a hold on the valence electrons