periodic trends + bonds and polarity

atomic radius

Distance from the nucleus to the valence shell

dependent and independent variable in AR graph

• Atomic # = independent variable 

• Atomic radius = dependent variable 

nuclear charge (Z)

•  positive charge present in the nucleus of an atom

• keeps electrons in place

shielding

• the repulsion of outer electrons by inner electrons 

• stronger than sideways repulsion

effective nuclear charge

• Net attraction of electrons to the nucleus 

• Z_eff = Z (nuclear charge) - shielding

sideways repulsion

• repulsion between electrons on the same shell

• causes spaces between electrons

• weaker than shielding

• mr.moniz term

trend of Z_eff

• across a period (left to right): protons increase while inner shielding electrons always stay the same, so effective nuclear charge also increases

• down a group: a new shell is added, all inner orbits become shielding electrons, increased shielding and low inner p+:e- causes more repulsion than attraction, so effective nuclear charge decreases

trend of atomic radius

• across a period (left to right): protons increasing, inner shielding electrons stay the same, so Z_eff increases, causes increased attraction, electrons pulled in more closely, atomic radius decreases

• down a group: AR does not increase because of an added orbit. A new valence orbit is added, inner orbits become shielding electrons, more repulsion, the p+:inner e- ratio decreases, valence is less drawn to the nucleus, atomic radius increases

ionization energy

• minimum amount of energy required to remove an electron from a neutral element in the gaseous state producing a +1 cation

• since minimum, it removes electron from valence shell

• X_g + IE = X_g⁺ + e- (free electron)

trend of ionization energy

• across periods (left to right): Z_eff increases, shielding stays the same while p+ increases, causing more attraction, more energy needed to remove e-, p+:inner e= ratio increasing, ionization energy increases

• down a group: Z_eff decreases, new shielding orbit, more repulsion than attraction, less energy required dislodge e-, p+:inner e- ratio decreases, ionization energy decreases

what does the spike in the ionization energy graph show?

• small spikes show jumps to inner orbit

• big spikes downwards show next row

• more energy needed for inner electrons due to increased effective nuclear charge

multiple ionization energies

• minimum energy required to remove subsequent electron to form a cation in an element in its gaseous state

• X_g⁺ + IE_2nd = X_g²⁺ + e- (+2 cation and free electron)

• change is discontinuous

• each subsequent electron needs more energy to be removed

why is more IE needed overtime? in one element, not as a trend

• Sideways repulsion is decreased because the first electron was removed from valence shell

• overall repulsion decreases so the attraction is more effective

• attraction itself does not increase but bc repulsion is less, overall attraction increases

electron affinity

• energy released by a neutral atom in the gaseous state when it attracts an electron to form a -1 anion (negative energy)

•  X_(g) + e^-  → X^-_(g) + E.A

why is energy releases when an electron is added to a neutral atom

• the electron is attracted to the positively charged nucleus

• falls into atom

• the potential energy is being converted into kinetic energy

• as it loses energy, that energy must be released

• energy released as light once electron reaches atom, this energy is EA

what does it mean if an electron needs to be pushed into an atom (EA)?

• that means there is repulsion (from the other electrons there)

• repulsion to add electron: negative electron affinity written as a positive number

electronegativity

a calculated average of measurements indicating the tendency of an element to attract an electron charge towards it in a compound

What factors allow metallic elements to form ions where they have lost their valence electrons?

• Metallic elements effective nuclear charge is low

• IE is also low

• The proton to inner electron ratio is lower

• so metallic elements are more likely to form cations because they have lower ionization energy

stable

how difficult is something to change, is it difficult or easy to change 

electron affinity

energy released by a neutral atom in the gaseous state when it attracts an electron to form a -1 anion, exothermic

electronegativity and electron affinity trend

• across a period (left to right): both increase, EN: more protons, unchanged shielding, Z_eff increases, more attraction, greater tendency to attract electron. EA: more energy given off when anion forms

• down a group: both decrease, EN: new shielding, more repulsion than attraction, tendency to gain electron is less. EA: greater repulsion, less energy released when anion is formed

• noble gasses do not attract electrons because the valence shell is full, and much more energy is needed to overcome the repulsion for an electron to go into the next shell

• noble gasses have high IE and are less likely to lose an electron

all trends general

• left to right: Z_eff increases, AR decreases, IE increases, EN and EA increases

• down a group: Z_eff decreases, AR increases, IE decreases, EN and EA decreases

covalent bond

a bond where two nuclei of nonmetals are bonded together by their mutual attraction to the same pair of electrons

how does a covalent bond occur?

• two nonmetal atoms collide so that the nuclei are closer than two times the radius

• the electron pair is drawn to the internuclear distance where they can create a concentration of charge

• side ways repulsion increased

• mutual attraction to same pair of electron bonds atoms together

• creates a covalent bond

differences between ionic and covalent bonds

ionic:

• non localized and non directional

• all ions pack together

• forms ionic crystal lattice

covalent

• attraction is localized to the internuclear distance

• attraction is directional along internuclear distance

• ex. in a H₂ bond, another extra hydrogen will not bond as no attraction is available

• forms molecule

How and why do ionic compounds form?

• a non metal and metal atom must physically collide with enough energy that an electron from the metal is dislodged

• the metal has low IE, easier to remove electron

• the non metal has high EN and EA, the dislodged electron is attracted to the non metal

it is always true that…

丨IE丨 (energy required) > 丨EA丨 (energy released) therefore the electron transfer from metal to nonmetal always requires energy and the ions are less stable than the neutral atoms

how does an ionic crystal lattice form?

• the positive and negative ions attract

• every cation surrounded by a number of anions and so on to maximize attraction

• a great amount of energy is required to create the crystal lattice, but much more is needed to break it apart, so it is stable

• ionic compound is more stable than neutral elements due to high attraction necessary to break apart

endothermic vs exothermic

• Endothermic: Process that absorbs/requires energy 

• Exothermic: releases energy 

Why does covalent bonding result in molecules and not ionic crystal lattices?

• ionic bond: metal can attract nonmetal as long as they are relatively close, it is non directional and non localized, which is why they pack together to create a crystal lattice

• covalent bond: the attraction is localized and directional to the internuclear distance and not just anywhere, so ions cannot pack together

what type of bond is this?

single bond

what type of bond is this?

double bond

polarity

uneven distribution of netgative charge in a bond

Partial charges: δ ⁺ or δ^-

• used to indicate the polarity of a covalent bond

• detla positive on less EN element

• detla negative on more EN element

how do you predict what kind of bond will form between elements in a compound?

use differences in electronegativity

electronegativity differences and type of bond

• 0≤ Δ EN≤ 0.4, non-polar or slightly polar

• 0.4 < Δ EN < 1.7, polar covalent bond

• Δ EN ≥ 1.7, ionic compound

radius of cations and anions compared to their neutral atom

• radius of cations are smaller than their neutral atom since there are less electrons, more overall attraction

• radius of anions are greater than their neutral atom since there are more electrons, more overall repulsion