Periodic Trends Flashcards
First ionization energy
- The amount of energy needed to remove the first electron from a neutral atom (the outermost electron)
- In general, first ionization energy increases going up and right in the periodic table
Effective nuclear charge and ionization energy
Elements with a higher effective nuclear charge have a higher first ionization energy because of the increased attraction from the protons making it more difficult to remove electrons from the atom (“ionize them”)
Ionization energy trend:
Increases going to the right of the periodic table and going up the periodic table (noble gases are included)
Elements with the _____ first ionization energy are the ones that will most readily lose an electron and become ionized
LOWEST
Ex: potassium has a lower ionization energy than lithium and therefore will most readily lose an electron
Ionization energy
The amount of energy required to remove an electron from an atom. The easier it is to remove an electron, the lower the ionization energy
General trend form atomic radius for atoms with a neutral charge
Increases going down and to the left of the periodic table
Atoms get larger going down: due to higher energy “n” shell = larger size
Atoms get smaller going across: due to higher effective nuclear charge = smaller size
General trend for atomic radius for isoelectronic ions
More positive charge = smaller radius
More protons = more attraction to electrons = smaller size atom
More negative charge = larger radius
Less protons = less attraction to electrons = larger size atom
What are the two factors that affect the size of an ion?
The number of atomic orbitals: an ion with more atomic orbitals will have a larger atomic radius because these atomic orbitals occupy more space and expand the electron cloud
The effective nuclear charge: an ion with fewer protons will have a larger atomic radius because protons pull in the electron cloud, reducing its range. The fewer protons in the nucleus, the weaker the pull on electrons
What is the reason for atoms size increasing going down a column on the periodic table?
This is due to electron shielding
The electron shielding effect reduces the effect of the nuclear charge. As more core electrons are added, they begin to shield the valence electrons from the increasingly positive nucleus which allows the outer electrons to move further away from the nucleus and therefore increasing the atoms size
Electron affinity
The amount of energy released when an atom gains an electron
Electronegativity
- The ability of an atom to attract electrons in a bond
- The higher the electronegativity of an atom, the greater ability it has to attract an electron pair
Which atoms would most readily form ionic compounds?
The ones with the lowest ionization energy (furthest to the left and down a column) because they will most readily lose electrons to become ionized
Noble gases exert what type of force?
Very weak london dispersion forces (and therefore have very low melting and boiling points)
Characteristics of halogens:
7 valence electrons
High electronegativity
High electron affinity
Easily reduced (gain one electron to have a full octet)
Strong oxidizing agents
Highly reactive with metals
properties of metals
- malleable, lustrous
- good conductors of heat/electricity
- form basic oxides
- lose electrons to form cations
- usually solid at room temperature (except Hg-liquid)
- high melting and boiling points
properties of non-metals
- brittle, dull
- poor conductors of electricity/heat
- form acidic oxides
- gain electrons to form anions
- gas or solid at room temperature (except Br-liquid)
- low melting and boiling points
trend when comparing melting points of compounds
comparing their ionic character: difference in their electronegative –> the larger the difference in electronegativity, the greater their ionic character and the higher the melting point
trend for metallic character:
increases going from right to left and going down a group
trend for electronegavity:
- increases going to the right and up the periodic table (excluding the noble gases
- F is the most EN
noble gases have full valence shells –> no EN value
what is the reasoning for the EN trend?
- increasing EN from left to right: increasing number of protons increases ability of an atom to attract an electron pair
- decreasing EN going down a group: as atomic radius increase, the valence electrons experience greater electron sheilding and the ability of an atom to attract an electron pair is decreased
electron affinity trend:
- increases going to the right and up on the periodic table (excludes noble gases)
reasoning for electron affinity trend:
- increas in electron affinity going to the right: greater nuclear attraction between protons and electrons causes a stronger affinity for electrons
- decrease in electron affinity going down a group: attraction of electrons to the nucleus decreases due to electron sheilding causing a decrease in electron affinity
what atoms are exceptions in the electron affinity trend?
- Phosphorus and silicon: even though P is to the right of S, P has a half-filled subshell electron configuration –> S has a higher electron affinity because gaining an electron will give it the half filled subshell configuration (more stable)
difference between electronegativity and electron affinity?
electronegativity is specific to electron pairs in a bond whereas electron affinity is the addition os a single electron
ionization energy of alkaline earth metals and group 13 exception:
- alkaline earth metals (group 2) have a greater ionization energy than group 13 elements
- alkaline earth metals have full filled orbitals and require more energy to remove an electron than group 13 metals
ionization energy of group 15 and group 16 exception:
- group 15 elements have higher ionization energy than group 16 elements
- group 15 have half-filled orbitals, which is a more stable configuration and requires more energy to remove an electron
ionic bonds
- complete transfer of electrons between 2 atoms
- metal loses electrons to become a positively charged cation
- non-metal accepts the electron and becomes a negatively charged anion
sigma bonds
- covalent bonds made by end-to-end overlap of atomic orbitals
- stronger than pi bonds
pi bonds
- covalent bonds made by side-to-side (lateral) overlap of p-orbitals
