unit 5: trends and bonding

Unit 5: Trends and Bonding

Valance electrons: available electrons to be lost, gained, or shared in the formation of compounds

Ionic Bond: losing or gaining electrons to form a bond between ions, creating a neutral compound

Ion: an atom or group of bonded atoms that has a positive or negative charge.

Cation: When metals lose electrons to form positive ions

Anion: When non-metals gain electrons to form negative ions.

Covalent Bond: sharing of electrons to form neutral compounds

Charges on Periodic Table+1

+2 +3

0

-1-2-3+/-4

The Periodic LawMendeleev: created the first periodic table to relate the

properties of elements and arranged them according to atomic mass

• Problems: (1) Most elements could be arranged in order of increasing atomic mass but a few could not?(2) What was the reason for chemical periodicity?

Mosely: arranged elements based on atomic numberPeriodic Law: the physical and chemical properties of the

elements are periodic functions of their atomic numberPeriodic Table: an arrangement of elements in order of

atomic number so that elements with similar properties fall into the same group or column.

s-Block Elements: Groups 1 &2• group 1 elements are more reactive than group 2

because it is easier to remove 1 electron rather than 2• Alkali metals: elements of group 1, highly reactive• Alkaline-Earth metals: elements of group 2, very

reactive

d-Block Elements: Group 3-12. • Transition elements that are typically less reactive than

s-block elements

p-block elements: Groups 13-18• properties vary greatly, metals, non-metals, metalloids• Halogens: group 17 elements, most reactive non-metals

– Ability to react is based on them having 7 electrons in outer shell, (they want 8 to be stable)

• Noble Gases: group 18 elements, stable and unreactive.– They already have 8 valence e-

f-block elements: Lanthanides and Actinides

Lanthanides: shiny metals similar in reactivity to group 2

Actinides: all but first four are man made in laboratories

(1) Atomic radius - one-half the distance between the nuclei of identical atoms that are bonded together.

r = d/2

• Period: AR decreases as you go across a period. (leaving noble gases out of it)– More e- in an energy level, the smaller it is,

because of the increased attraction of the negative to the positive nucleus.

• Group: AR increases as you go down a group. - adding extra layers of electrons- more energy levels

Trends in Periodic Table

Atomic RadiusS

M

L

(2) Ionic Radius: Ion: an atom that has a positive or negative chargeCation: a positive ion, from loss of electrons

– decreases atomic radius b/c less electrons (energy level)

Anion: negative ion, from addition of electrons– increase atomic radius, increases electron repulsion

• Metals tend to form cations• Non-metals tend to form anions

– Period: decreases across a period (like atomic radius)

– Group: increases down a group (like atomic radius)

Ionic Radius

SM

L

(3)Ionization Energy - amount of energy required to remove the outermost electron from a neutral atom -to make positive ions

The energy you need to put into an atom to take away an electron. (give and take) They give you an e- because you gave them energy

• Ionization: the formation of an ion• Period: IE increases as you go across a period

– When elements are closer to having a full octet, they do not want to give up an electron, so IE is much higher

• Group: IE decreases as you go down a group. Or increases as you go up a group– it is easier to remove the electrons from outer energy

levels because they are farther from the positive nucleus and thus less pull by the nucleus

Ionization Energy

S

M L

(4) Electron Affinity: the energy change that occurs when an electron is acquired by a neutral atom-to make negative ions

• The energy given/ released when you add an electron to an atom (give and take)

• They give you energy because you give them an e-• “love of electrons”• Most atoms release energy when they acquire an

electron, A + e- A- + energy• Period: increases as you go across, but it is negative

energy because the atom is releasing it. The more they want electrons the more they will “pay” for it in energy, thus giving more energy

• Group: decrease as you go down because electrons add with greater difficulty going down a group, because of increased atomic radius. They don’t want e- so they won’t “pay” much for it.

• Non-metals: Some atoms want electrons, so it is easy to give them one and they are grateful and give you energy in return. So they release energy (negative E)

• Metals: Some atoms do not want electrons, so you have to force the electron onto the atom by also using/giving energy. So they absorb energy (positive E)

Electron Affinity

S

M L

(5) Electronegativity: an atom's ability to grab another atom's electrons (ability to attract electrons)

• Occurs in covalent bonding, when electrons are shared• Period - increases as you go across a period

– Because atoms want electrons more as you go across

• Group - decreases as you go down a group or remain the same. – Since there are more energy levels the positively

charged nucleus cannot attract the electrons well enough

Electronegativity4.0

0.7

Electronegativity

(6) Reactivity - refers to how likely an atom is to react with other substancesMetals: the more likely to lose an electron…the more reactiveNon-metals: the more likely to gain an electron….the more reactive.

Reactivity

L

L

Non-metals

Metals

Melting Point and Boiling Point• Metals: Decreases as you go down a group

and increases as you go across a period

• Non-metals: Increases as you go down a group. Generally decreases as you go across a period

S

S

M L

L

Electron Dot Notation: electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol.

Remember- orbitals fill with one electron before any one fills with two electron (Hund’s Rule)

Octet Rule: compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has a full octet (8 e-) in its highest occupied energy level

When electrons are shared, orbitals are overlapped.• To form a bond: release energy • To break a bond: absorb energy, it requires energy

BONDING:

Electron-dot can be used to represent molecules:

Unshared Pair (lone pair): pair of electrons that is not involved in bonding and that belongs exclusively to one atom. (not bonded)

Lewis Structures: ( F F )– dot-pairs or dashes between two atomic symbols represent

electron pairs in covalent bonds

– dots adjacent to only one atomic symbol represent unshared electrons

Two Types of Compounds• There are two ways to achieve a full octet to become

stable- Ionic bonding or Covalent Bonding

1) Ionic compounds: result from ionic bonding• Metals and Non-metals combine• composed of positive and negative ions that are

combined so that the numbers of positive and negative charges are equal

• group 1or 2 want to give electrons away and group 16 or 17 want to take electrons, so valence electrons are transferred between atoms upon collision.

• All metals can combine with a non-metal to form Ionic bonds through electron transfer.

• Show the transfer with charges!• Positive ions- lose electrons• Negative ions- gain electrons

• Forces that hold ions together is very strong

• Ca (+2) and F (-1) give you CaF2 (1-to-2 ratio)

Electron-Dot: Ca F2:

Na2O:

What are Ionic compounds like?: • salt, like ocean water.• Strong and brittle like egg shells, once you crack the

shell it all falls to pieces. • 3-D, tightly bound structure, crystal lattice structure

• most ionic compounds exist as crystalline solids• Lattice energy: the energy released when one mol of a

crystalline/ionic compound is formed• Ionic bonds are stronger than covalent bonds.

Crystal Lattice

2) Covalent Compounds: when electrons are shared• Bonding between Non-metals and Non-metals• when neither atom wants to give up electrons fully• Hydrogen always forms covalent bonds because a

Hydrogen can’t lose its only electron by transfer so it must share– H2O:

• covalent bond: atoms share electrons (no ions)• molecule: neutral group of atoms held together by

covalent bonds• Diatomic molecule: bond between 2 atoms of 1 element.

What are Covalent Compounds like?:• In a purely covalent bond, the shared electrons are

“owned” equally, no difference in electronegativity • Bonds within covalent compounds are strong, but

attraction between molecules are weak, so they will break apart into molecules easily.

• Small difference in electronegativity, so they must be close together on the periodic chart.

• Electrons are shared, so don’t write charges!– CO2

– [NH4]+1

• More than two electrons can be shared CO2

•In covalent bonding the initial repulsion between like subatomic particles, is overcome by the attraction between the positive nucleus and negative electron cloud, forming a bond of shared electrons.

Higher Potential EnergyUnstable

Lower Potential EnergyStable - Neutral

Steps for Covalent Structures1) Cross structure (carbon or least electronegative

atoms is in the middle-symmetrical)2) Give them what they want (8 e-, except H)3) Count what they want4) Determine how many valence electrons they will

bring to the compound (allowed)5) Remove the difference in pairs of e-, then move

a pair from the atom next to it, to be shared between them

6) Check- count total, and make sure that none have more than 8

• Ionic: MgCl2

• Covalent: CH2S Want: 14e-C: 4e-H:1e- x 2= 2e-S: 6e-Allowed: 12e-

Exceptions to Octet Rule:• Boron- only wants 6 e- in outer shell - BF3

• Beryllium- only wants 4 e- in outer shell- BeF2

• Expanded Octet/ Expanded Valance: elements that want more than 8 e- in outer shell, when they combine with highly electronegative elements (F, O, Cl)– The use the s (2), the p (6), and some of the d block

electrons

– PCl5 and SF6

Single Bond: covalent bond produced by the sharing of one pair of electrons between two atoms.

CH3I

Double Bond: covalent bond produced by sharing of two pairs of electrons between two atoms.

Triple Bond: covalent bond produced by sharing of three pairs of electrons between two atoms.

Multiple Bonds: double or triple bonds. • Bond Strength: Triple > Double > Single • Bond Energies: Triple > Double > Single• Bond Length: Single > Double > Triple• So Triple bonds are strongest, so they require

more energy to break, but they are the shortest.• Bond Length- the average distance between

two atoms• Bond Energy- the energy required to break a

bond to form single isolated atoms

Bond Energies and Heats of Reaction (H)Bond Energy is the energy required to break a chemical bond. Tabulated values (Table 6-1) are average bond energies in units

of kJ / mole. Bond-breaking is endothermic, bond-making is exothermic. H for a reaction can be estimated from bond energies as follows.

(Counting ALL bond energies as positive values!) H BE (bonds broken) - BE (bonds formed) Problem: Use data in Table 6-1 to estimate H° for the reaction.

CH2=CH2 + H2O CH3-CH2-OHBonds Broken Bonds FormedC=C 612 C-C 348H-O 463 C-H 412 = 1,075 C-O 360

= 1,120H° 1,075 - 1,120 - 45 kJ/mole

Resonance Structures:

• Some molecules cannot be represented by a single Lewis Structure. One such molecule is Ozone (O3)

• This is called Resonance and the different Lewis structures are called resonance structures. To indicate resonance a double-headed arrow is placed in between

Polyatomic Ions• A group of atoms that are covalently bonded but result

in an imbalance of electrons, thus creating an overall charge

• These ions are involved in ionic bonding, but also have covalent bonds within the ion involved

• The covalent cross structure has too many electrons- then a negative charge for the entire group is given.– Nitrate: NO3

- excess of electrons

• The covalent cross structure does not have enough electrons- then a positive charge for the entire group is given– Ammonium: NH4

+ shortage of electrons – their missing!

Polar Covalent: have partial charges

+ represents a partial positive charge

- represents a partial negative charge

HCl: Hydrogen and Chlorine. • Chlorine is more electronegative so when sharing

electrons is attracts electrons more, so it is partially negative

• Hydrogen wants electrons less so partial positive

Non Polar Covalent : two elements with similar desire for electrons so has NO partial charges, because of equal sharing, usually gases

Overall…..• Like dissolves Like: so charged ions and molecules will

dissolve each other and not those without charges. • Ionic- give and take- transfer of e-, thus full charges

• Polar Covalent- must share, but desire for e- is slightly different, so unequal sharing and partial charges

• Non-polar Covalent- same desire for e-, so equal sharing and no charges

Ionic vs. Covalent

Elements combine to form either ions or molecules. Properties of Ionic Compounds:

- Physical: strong, hard, brittle, well organized, tightly bound, 3-D crystal structures- takes a lot of energy to break bonds- tend to dissolve in water (like salt)- electrolytes: any compound that conducts electricity- large difference in electronegativity: greater than 1.7- strong Ionic character (greater than 50%)- stronger compounds- higher melting points- Ex: salt, egg shells

Ionic vs. CovalentProperties of Covalent Compounds:

- Molecules held together by strong covalent bonds- strong bonds within molecules, but a weak bond that holds one molecule to another.- Polar : liquids or soft solids that don’t conduct electricity (sugar)

- small difference in electronegativity: 0.3-1.7 difference - weak compounds

- low melting points

-Non-Polar: gases that don’t conduct electricity- almost no difference in electronegativity: < 0.3 difference- very weak compounds- lowest melting points

-Ex: Candles, plastics, crayons, diamonds

Molecular Geometry• Properties of molecules depend on bonding & geometry

Molecular polarity: uneven distribution of molecular charges

VSEPR Theory: (used to predict geometry) states that repulsion between valance electrons surrounding an atom causes them to be oriented as far apart as possible.

Bond Angles : AB2 (2 or 3 atoms) AB3 (4 atoms) AB4 (5 atoms)

Linear: - AB or AB2

- Example: HCl or CO2

- 0 lone pairs of e-- 2 atoms bonded to the central atom

Bent or Angular: - AB2(E)- Example: SnCl2

- 1 lone pair of e-- 2 atoms bonded to the central atom

Trigonal Planar: - AB3

- Example: BF3


Tetrahedral: - AB4

- Example: CH4


Trigonal Pyramidal: - AB3(E)- Example: NH3


Bent or Angular: - AB2(E2)

- Example: H2O - 2 lone pair of e-

- 2 atoms bonded to the central atom

Trigonal Bipyramidal: - AB5

- Example: PCl5


Octehedral: - AB6

- Example: SF6


Central Atom with 2 or 3 entities surrounding: 180o and 120o

Central Atom with 4 entities surrounding : 109.5o

Trigonal

Central Atom with 5 entities

surrounding

:90o and 120o

Central Atom with 6 entities surrounding: 90o and 90o

Hybridization• The mixing of two different orbitals

(blocks) to form a new set of orbitals

• The new orbitals have an energy somewhere between the original energies.

HybridizationHybridization: mixture of two or more atomic orbitals

of similar energies on the same atom to produce new orbitals of equal energies

Ex: Methane __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 1s sp3

Hybrid Orbitals: orbitals of equal energy produced by the combination of two or more orbitals on the same atom

sp - linearsp2 - trigonal planarsp3 - tetrahedral

Intermolecular forces: forces of attraction between molecules (covalently bonded compound)– These forces are weaker than bonds– The stronger the bonds, the higher the boiling point, the

stronger the intermolecular forces.Polar molecules have the strongest intermolecular forces

– polar molecules act as dipoles – Each molecule has an uneven partial charge…creating a

dipole.

Dipole: created by equal but opposite charges that are separated by a short distance. It comes from unequal sharing of electrons in polar molecules

(H Cl)

Intermolecular Forces

Dipole-dipole forces: forces of attraction between two polar molecules. (like in water)

Hydrogen Bonding:a hydrogen atom that’s bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule– Strong dipole-dipole forces, intermolecular forces– NOT a BOND, but a FORCE

London Dispersion Forces: intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles– even nonpolar molecules experience weak

intermolecular attractions

Dipole-Dipole Forces• Many molecules contain bonds that fall between the

extremes of ionic and covalent bonds. • The difference between the electronegativities of the

atoms in these molecules is large enough that the electrons aren't shared equally, and yet small enough that the electrons aren't drawn exclusively to one of the atoms to form positive and negative ions.

• The bonds in these molecules are said to be polar, because they have positive and negative ends, or poles, and the molecules are often said to have a dipole moment.

• HCl molecules, for example, have a dipole moment because the hydrogen atom has a slight positive charge and the chlorine atom has a slight negative charge.

• Because of the force of attraction between oppositely charged particles, there is a small dipole-dipole force of attraction between adjacent HCl molecules.

• These forces are week compared to ionic compounds, but they are the strongest among covalently bonded compounds

Dipole-Dipole Forces

Dipole-Induced Dipole Forces

• A polar molecule can induce a dipole in a non-polar molecule by temporarily attracting its electrons

• What would happen if we mixed HCl with argon, which has no dipole moment? The electrons on an argon atom are distributed homogeneously around the nucleus of the atom. But these electrons are in constant motion.

• When an argon atom comes close to a polar HCl molecule, the electrons can shift to one side of the nucleus to produce a very small dipole moment that lasts for only an instant.

Dipole-Induced Dipole Forces

• By distorting the distribution of electrons around the argon atom, the polar HCl molecule induces a small dipole moment on this atom, which creates a weak dipole-induced dipole force of attraction between the HCl molecule and the Ar atom.

• This force is weak.

Induced Dipole-Induced Dipole Forces

• By itself, a helium atom is perfectly symmetrical. But movement of the electrons around the nuclei of a pair of neighboring helium atoms can become synchronized so that each atom simultaneously obtains an induced dipole moment.

• London Dispersion Forces – instantaneous and short lived weak attraction when electron distribution is slightly uneven (caused by random movement of e-)

• Occur in Noble Gases and Non-polar molecues

Induced Dipole-Induced Dipole Forces

• These fluctuations in electron density occur constantly, creating an induced dipole-induced dipole force of attraction between pairs of atoms.

• As might be expected, this force is very weak and lasts for only a moment.

• Larger molecules are at particular risk of this because they have more electrons farther from the nucleus moving around, thus more likely to caused a momentary shift.

unit 5: trends and bonding

Documents

elements of group

properties of elements

atomic radius group

arrangement of elements

periodwhen elements

transition elements

reactivedblock elements

atomic massproblems