An Introduction to the Periodic Table
Column/Families/Groups
- Tell us how many valence electrons are in the outermost s and p shells
- Elements with similar properties are found in the same group
Rows/Periods
- Tell us which shell the valence electrons are found; each row has valence electrons in same energy lvl
Groups of Periodic Table Group 1: Alkali Metals
- Low density, soft, silver, very reactive, all salts, form strong bonds with water
Group 2: Alkali Earth Metals
- Stronger, denser, less reactive, one more valence electron
Group 3-12: Transition metals
- D-block elements Diagonal row: Metalloids
- Poor conductors, both properties Next: Nonmetals
- Nonconductors, dull, brittle
Group 7: Halogens
- Most reactive cuz have 7 ve
Group 8: Noble Gasses
- Non reactive cuz have 8 valence electrons and are stable
○ 2 in the s and 6 in the p
- Most gasses at room temperature
- When solid take crystal form
- Can run electricity thru noble gasses but won’t react
Polyelectronic Atoms
- Polyelectronic atoms: atoms with more than one electron
- Three energy contributions in the description of an atom
○ The kinetic energy of the electrons as they move around the nucleus
○ Effective Nuclear Charge
○ The potential energy of repulsion between electrons.
Effective Nuclear Charge
● Effective nuclear charge (Zeff): attraction between the nucleus and the valence electrons
○ Note: nuclear charge is the total (+) charge of the protons and attraction to all the
○ 2nd Note: Nuclear charge increases both down a group and across a period; Zeff weakens down a group and increases across a period
■ Zeff depends more on distance than # of protons
● When justifying trends talk about Zeff and why
Weakens moving down a group
● Attraction between valence electrons and nucleus weakens bcuz of increasing number of filled energy lvls and thus distance between them
- Valence electrons are shielded more from the nucleus
- Coulomb’s law: attraction between 2 charged particles is inversely related to the distance
Strengthens from left to right across a period
● Electrons are being added to the same principal energy lvl & are more strongly attracted to the nucleus due to additional protons
- Valence electrons are less shielded from the nucleus (due to added e- pair repulsions)
- Coulomb’s law: attraction between 2 charged particles in directly related to the magnitude of their charges
Atomic Radius
● Size of an atom, distance from the nucleus to the outermost electrons (half the distance between two nuclei of a molecule)
● Is determined by how much the electrons are attracted to the (+) nucleus
- Number of protons and electrons determines the size of atoms and ions
Atomic Radius Trends
● Group Trend: increases as you go down
- Bcuz Zeff decreases (less attraction = larger size)
● Periodic Trend: decreases as you go across (From left to right)
- Bcuz Zeff increases → (more (+) nuclear charge = more attraction= smaller size)
Ionic Radius
● Size of an atom when it is an ion
Ionic Radius Trend
● Metals lose electrons → highest occupied energy shell of atom is further away from highest occupied energy shell in ion (Stronger Zeff so…)
- Ionic radius<atomic radius
● Nonmetals gain → more e- that p+ (weaker Zeff bcuz of stronger electron-electron repulsions so…)
- Ionic radius>atomic radius
● Group Trend: increases as you go down
● Period Trend: decreases as you go across
● Explain Trends
- Big jump between Carbon and nitrogen ions (10 e-) bcuz they have different number of energy lvls
- Decrease nitrogen to fluorine bcuz increasing Zeff
Isoelectronic Ions
● Ions containing the same number of total electrons:
- All have 10 electrons and the electron configuration
Ionization Energy
- Energy needed to remove a valence electron and form an ion
- First ionization energy (I1): the energy required to remove the highest-energy (outermost) electron of an atom
○ The amount of energy required to remove a second or third electron from an atom is always more than the amount of energy required to remove the first electron
■ I1 < I2 < I3
○ Why? Bcuz the Zeff increases as you remove electrons bcuz of reduced electron-pair repulsions & e- are closer to nucleus→ attraction that remaining electrons feel is greater
- Very important when explaining ionization energy trends
○ It takes more energy to remove core electrons than valence e- (in outermost s and p)
○ Takes more energy to remove an electron from a half or fully filled electron shell
- Exceptions in Ionization Energy trend
○ Slight decrease in IE from beryllium to boron (group 2 to group 13) bcuz more energy is required to remove an electron from a filled 3s sublvl than from a partially filled 3p sublvl
○ Slight decrease in IE from from nitrogen to oxygen (group 15 to 16) bcuz it takes more energy to remove an electron from a half filled 2p sublvl than from a partially filled 2p sublvl (due to the added electron repulsions)
- Ex: Explain this trend for Mg
○ Have to answer…
■ a) basic reason for trend (see above)
■ b) why such a big jump between I2 and I3
- Bcuz after removing I2 you are no longer removing valence electrons but core electrons from a filled energy lvl → takes much more energy
- Note: AP exam often focuses on this trend
Ionization Energy Trend
- Group Trend: decreases down a group
○ Because Zeff is weakening
- Periodic Trend: increases across a period
○ Because Zeff strengthens
Electronegativity
- Measure of how much an atom in a molecule attracts shared electrons to itself; affected by atomic
radius and atomic number
○ Associated with the production of a negative ion Electronegativity Trend
- Group Trend: decreases as you go down
○ Zeff decreases
- Period Trend: increases across
○ Zeff increases; atom less likely to give up electron
Electron Affinity
- The energy change associated with adding an electron to a gaseous atom
○ The more negative the energy, the more energy is released (exothermic)
Metals
- Metals tend to have positive electron affinities while nonmetals tend to have negative electron
affinities
Reactivity
- “Ability of an atom or compound to undergo a chemical reaction with another atom”
- Metals lose e- when react so reactivity based on low ionization energy
○ Low I.E = high reactivity
- Nonmetals gain e- when react so reactivity based on high electronegativity
○ High electronegativity = high reactivity
Properties of Nonmetals
- The ability to gain one or more electrons to form an anion when reacting with a metal. Thus nonmetals are elements with large ionization energies and the most negative electron affinities.
- Non-conductors (no flow of electrons)
- Dull, brittle
- Most gasses at room temperature and form crystals when solid Metalloids
- exhibit both metallic and nonmetallic properties under certain circumstances
The Properties of a Group: The Alkali Metals
Alkali Metals
- Most reactive of metals, low density, soft, silver, all salts, form strong bonds with water
- Note: even though hydrogen is in Group 1A of the periodic table, it behaves as a nonmetal, bcuz of its very small size → electron in the small 1s orbital is bound tightly to the nucleus