Chem Unit 11: The Periodic Table & Bonding

2023

Atomic Radius Definition

½ of the distance between the nuclei of 2 atoms of the same element when the ats are joined

Atomic Radius group trend

Atomic size ^^increases^^ from top to bottom within a group

Why (AR group trend)

%%As atomic # increases within a group, charge on nucleus increases(more p+ and more e-), # occupied energy levels increase.%% Shielding effect greater than the effect of the increase in nuclear charge, therefore atomic size decreases.

AR Period trend

Atomic size ^^decreases^^ left to right across a period

Why? (AR Period trend)

Atomic size decreases across Period left to right. Each element has 1 more e- than preceding. e- added to same principal energy level, shielding effect constant for all elements in a period. $Increase nuclear charge pulls e- in highest energy level to nucleus, atomic size decreases.$

Period

Left to right

Group

Top to bottom

Largest element (atomic radius) located

Bottom left

Smallest element (atomic radius) located

Top right

Ionization energy definition

The energy required to remove an e- from an atom

When is it the hardest to remove e- (ionization energy)

Full energy level, therefore ==easier to take from metals==

Group trend (IE)

1st ionization energy ^^decrease^^ from top to bottom

1st ionization energy def

Energy required to remove 1st e- from an atom

Why?( group trend IE)

The atomiz size increases as atomic # increases within a group. As size of atom increase, nuclear charge has a small effect on the e- in highest occupied E.L., less energy required remove e- from this energy level

Period trend (IE)

^^Increases^^ left to right

Why?(period trend IE)

As nuclear charge increases, shielding effect stays constant. Nuclear charge increases across a period, increase attraction of nucleus for an e-, it takes more energy to remove an e- from an atom

Electronegativity def

%%Ability of an atom of an element to attract e-%% (when the atom is in a compound)

$Nonmetallic better at$

Group trend (EN)

Top to bottom ^^decreases^^

Dmitri Mendeleev & PT

\-invented PT (1869)

\-$arranged elements by atomic mass$

\-predicted existence unknown elements

\-element Md

Henry Moseley & PT

\-determined atomic # of elements from theirnX-ray spectra (1914, worked with Rutherford)

@@-arranged elements by increasing atomic #@@

Periodic Law

When elements are arranged in order of increasing atomic #, there is a periodic pattern in their chemical and physical properties

Physical properties (examples)

Color, luster, malleability, density, MP, BP, FP, conductivity

Chemical properties (examples)

Chemical reactivity, flammability, reactivity w/ acids, oxidation

Metals

80%, good conductors, solid @ RT (except Hg)

Metalloids (semi-metals)

“Steps”, depends on conditions

Non-metals

Top right, most gas @ RT, more variation, poor conductors (except C)

Representative elements (regional/group name)

Groups 1, 2 , 13-17 or “A”

Transition elements

Group 3-12 or “B”

Inner transition elements

“Rare-earth” elements

Halogens

Group 17, non-metal

Noble gases

Group 18, “inert" gases- not very reactive

Why?(group trend EN)

Metals far left of PR have low values, nonmetallic far right have high values

Period trend EN

Left to right %%increases%%

Purpose of bonding

To fill the valence shell of an atom

Most elements are stable when

They have the same # of valence e- as a noble gas (8)

Atoms will__electrons in order to become stable

Lose, gain, or share

Chemical bond

Mutual attraction of 2 molecules on an e-

When a chemical bond is formed, energy is

Released, exothermic

When energy released as a bond formed, the resulting system is

At lower energy, more stable

When chemical bond is broken, energy is

Absorbed, endothermic

How much energy is needed to break the bond that was released when the original body was formed

Same amount

If little energy is released in the formation of a bond, it requires only

A small amount of energy to break it, said to be a stable bond

Four kinds of bonds:

Ionic, metallic, network, covalent

Ionic bonds

A $bond between a metal and a non-metal$ made by the creation of ions that are %%electrically opposite and therefore attracted to each other%%

Cation + anion =

Ionic crystal (neutral)

Ionic bonds energetically favorable to get __ e-

8

Atoms react in order to achieve

The stable electeon configuration of a noble gas

Cations =

* Metals
* Formed from loss of e- (oxidation)
* Creates a + charged particl

Anions=

* Non-metals
* Formed from gain of e- (reduction)
* Creates - charged particle

Ionic bond formation

* Electromagnetic attraction of oppositely charged ions
* Transfer e- from metal to non-metal
* Creates a formula unit (empirical formula)

Formula unit definition

Smallest sample (particle) of an ionic compound (like NaCl)

Ionic properties- ionic compounds tend to be:

* Crystalline
* Solid at room temp
* Have high melting and boiling points (usually >300°C)
* @@Conduct electricity when in the aqueous or liquid state(or gas)@@

Covalent bonds formed

When 2 nuclei ==share== e- equally or unequally

Covalent bons generally occur

Between 2 non-metals (EN's are so high neither atom would want to lose e-s)

Covalent bonds can form

* Single bonds (sharing 1 pair of e-)
* Double bonds (share 2 pairs)
* Triple bonds (share 3 pairs)

Types of covalent bonds

* Non-polar (non-dif) covalent bonds
* Polar covalent bonds

Non-polar covalent bonds formed

When 2 or more of th%%e same kind of atom%% are bonded together by sharing electrons, ex diatomics

Non-polar covalent bonds, since all atoms have the same attraction for electrons

(Same EN), the electrons are shared equally

Non-polar covalent bonds EN difference

0

N-P result of sharing e-

Is a molecule where each of the atoms has a full valence shell (usually 8 valence e-)

Polar covalent bonds- EN values

Unequal (EN difference does not = 0; usually >.4)

Polar covalent bonds- sharing of electrons

Unequal

EN difference of Polar covalent bonds

>.4

Partial charges Polar covalent bonds

(δ-) charge = \~ -

(δ+) charge = \~ +

Rules for drawing dor diagrams of Polar covalent bonds

Lewis dot diagrams do not indicate

Valence-Shell Electron-Pair Repulsion Theory (VSEPR)

* One method used to predict the geometry of specific molecules
* Atoms are bonded together in molecules by the shared pairs of e-
* Best arrangement of a molecule is one which minimizes repulsion of electron pairs

Bonding pairs e-

Shared between 2 atoms

Non-bonding e-

Not shared

Electron pair geometry

\

Molecular geometry

* Based on shared pairs of e- within a molecule
* When describing the shape of a molecule, we give the molecular geometry rather than the electron pair geometry
* The molecular geometry can be determined from the electron pair geometry

Steps in predicting the molecular geometry

Linear molecular geometry

2 total e- domains, electron-domain geometry = linear, bonding domains = 2, NB domains =0

Trigonal planar molecular geometry

3 pairs e- domains, trigonal planar electron- domain geometry, 3 bonding, 0 NB,

Bent molecular geometry (3 pairs)

3 pairs e- domains, trigonal planar electron- domain geometry, 2 bonding, 1 NB

Tetrahedral molecular geometry

4 pairs e- domains, tetrahedral electron- domain geometry, 4 bonding, 0 NB

Trigonal pyramidal molecular geometry

4 pairs e- domains, tetrahedral electron- domain geometry, 3 bonding, 1 NB

Bent molecular geometry (4 pairs)

4 pairs e- domains, tetrahedral electron- domain geometry, 2 bonding, 2 NB

Shapes chart

The bond angles decrease as

\# of non-bonding e- pairs increases

Effect of non-bonding on bond angles-reason:

* Bondinding pairs of e- are attached to both nuclei of the bonded atoms (are locked, in a bond)
* Non-bonding e- are not confined in a bond so they can move under the attractive influence of a nucleus → take up more space
* non-bonding e- will repel each other somewhat sausing the bonding pairs to move closer together

Effect of multiple bonds on bond angles

* Multiple bonds contain higher distribution charge (due to more e- pairs) than single bonds→ multiple bonds affect bond angles
* E- in multiple bonds will exert a greater repulsive force on adjacent e- pair than single bonds

Molecules with no single central atom

* Cannot describe this molecule as 1 type of molecular geometry

Type of molecule determined from

The distribution of charge

Polar =

Asymmetrical (uneven)

Non-polar =

Symmetrical (even)

Polar molecules produced when

There is an uneven distribution of electrical charge (dipole)

Conditions where Polar molecules exist

1. A molecule composed of %%only 3 elements with a Polar covalent bond%% is %%always%% polar
2. A molecule composed of $more than one non-metal atoms,$ bonded by Polar bonds, will always be Polar if the shape of the $molecule is asymmetrical$

Conditions where non-polar exist

Contains even distribution of charge (arrows cancel)


1. @@Diatomic@@ molecules that are bonded by non-polar bonds
2. ==>2 non-metal atoms, bonded by Polar covalent bonds, symmetrical shape==

Polar substance will dissolve in

Polar molecules

(Partial charge on Polar molecules are attracted to charges on Polar substance)

Ionic compounds dissolve in

Polar molecules

(Ionic charges (cations, anions) are attracted to Polar molecules)

Non-polar molecules will dissolve in

Non-polar molecules

Bonds between 2 atoms represetn

The force of attraction between atoms

H-Cl

Molecular attraction (IMF's)- physical properties s,l,g largely due to

* Molecular attraction
* Stronger = solid
* Weaker = gas

\

Why strength of attraction impact on phase?

Stronger the molecular attraction the higher the temp at which a substance melts ot boils (takes more energy to overcome the force of attraction)

Types of IMF

1. Dipole-Dipole
2. Hydrogen bonding attractions
3. Molecule-ion Attractions
4. VanderWaals Forces (London Dispersion Forces)

Dipole-Dipole attraction

* Exists between Polar molecules (uneven distribution of charge)
* Effective only when Polar molecules are very close together
* Weak attraction

H bonding attraction

* Special type dipole-dipole bonding that exists between molecules that contain H bonded to a small highly EN atom such as F,O, or N
* H only has a small share of electrons (EN difference is so high that electrons are pulled closer to O,F,N)
* Results in highly polarized bond and molecule (partial charges are stronger)
* Attractions between these highly polarized molecules is like an extra strong dipole-dipole
* H20, HF, NH3

Metallic bonds

* Metallic atoms have few valence e-
* Have low IE
* Metallic bond= collection of cations (+) for which free-floating valence e- have an attraction
* “A sea of mobile electrons”