Basic Chemistry 04

8/14/2019 Basic Chemistry 04

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CH 1200 UNIT 6 PERIODIC TABLE


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CH 1200 UNIT 6 PERIODIC TABLECollege of the North Atlantic Qatar

6.0 THE PERIODIC TABLE What is the periodic table ?What is the periodic table ?

What information is obtained from the table ?What information is obtained from the table ? How can elemental properties be predicted basedHow can elemental properties be predicted based

the periodic table ?the periodic table ?


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DimitriMendeleev(1869)

Dm it ri M ende lev, a

Russian chem ist , t ea cherand writer is considere dto be t he father of t hem odern periodic t able of

the e lement s

http://www.chem.msu.su/eng/misc/mendeleev/welcome.html


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6.0 The Periodic Table


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6.1 DESCRIPTION OF THE PERIODICTABLE

Periodic Law : When elements are

arranged according to Atomic Number,similar physical and chemicalproperties recur in a periodic fashion.

Period: A period consists of anyhorizontal row of the periodic table. Atpresent there are 7 periods on thePeriodic Table.

Group (Family): A group consists of anyvertical column of the Periodic Table.The more important Families havespecial names.


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6.1.2 PERIODS ON THE PERIODIC TABLE

There are 7 periods on the P.T.

Each horizontal row represents aperiod.

The first element of any period is ametallic element

The last element of any period is aNoble gas element.

As we move across each period, thecharacteristics change gradually frommetallic to non-metallic.


7/99CH 1200 UNIT 6 PERIODIC TABLECollege of the North Atlantic Qatar

6.1.2 Periods or Rows

Periods: Are horizontal rows across the

periodic table (rows 1-7).They have similar properties and

generally show trends.1

IA

18

VIIIA

12

IIA

13

IIIA

14

IVA

15

VA

16

VIA

17

VIIA

2

33

IIIB

4

IVB

5

VB

6

VIB

7

VIIB

8 9

VIIIB

10 11

IB

12

IIB

4

5

6

7

2nd Period

6th Period



6.1.3 FAMILIES ON THE PERIODIC TABLE

There are 8 important families on the

P.T. Each family can be identified by the

Number at the top of the familycolumn: IA , IIA ect.

Ex: Group IIA is the Alkaline EarthFamily Group VIIA are the Halogens Group VIII8 are the Noble Gases

Each family can be identified by any ofthe numbers from 1 18 which areshown at the top of each family.

Ex: Group 15 is the Nitrogen Family Group 16 is the Oxygen Family.



6.1.3 Families and Groups

The groups shown below share similar

properties


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6.1.3 - Families or Groups

Families: Are vertical columns down the

periodic table (columns or group, 1- 18)and generally show trends.These elements have the same number ofelectrons in the outer most shells, thevalence shell.1

IA

18

VIIIA

1 2IIA

13IIIA

14IVA

15VA

16VIA

17VIIA

2

33

IIIB

4

IVB

5

VB

6

VIB

7

VIIB

8 9

VIIIB

10 11

IB

12

IIB

4

5

6

7

Alkali Family:1 e- in the valence shell

Alkali Family:1 e- in the valence shell

Halogen Family:7 e- in the valenceshell

Halogen Family:7 e- in the valenceshell


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. - ren s n er o cTable

The elements of the

periodic table are arranged ina pattern.

These patterns show

trends in properties such as: Atomic Radius

Ionic Radius

Ionization Energy Electron Affinity

Electronegativity


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. . - ren n om cRadius

Note the trend in atomic radius shownbelow

The atomic radius increases down thegroup and also increases from rightto left across a period

6 2 1


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6.2.1 - ren n om cRadius

Atomic Radius increases down each group Atomic Radius increases from right to left

across period


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6.2.1 Trends in Atomic Radius


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6.2.1 Trends in Atomic Radius


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6.2.2 IONIC RADIUS

Ionic Radius is defined as the radius of

an ion.

Cation Formation usually produces anion which is smaller than the originalatom.

Anion formation usually produces anion which is larger than the originalatom.


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6.2.3 Atomic Radius vs Ionic Radius

ompar ng ren s n


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. . - ompar ng ren s n .T.

Moving Left --> Right

Atomic Radius Decreases

Ionization Energy Increases

Electronegativity Increases

Moving Top --> Bottom

Atomic Radius Increases Ionization Energy Decreases

Electronegativity Decreases

4 Tren n Ion zat on


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. .4 -Tren n Ion zat onEnergy

Ionizat ion Energy ( IE) : the quant it y of energy

required t o remove an elect ron from an a t om .The energy required to remove the valenceelectron(outer electron) from an atom. IE is largest toward top leftcorner of periodic table since these atoms hold on to their

valence e- the tightest.


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6.2.4 INOIZATION ENERGY VALUES

ren n on za on


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. . - ren n on za onEnergy

Ionization energy increases from left to

right and from bottom to top of theperiodic table


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6.2.5 Electron Affinity

Electron Affinity is defined as the

energy change that occurs when anelectron is added to an atom to forman anion.

An anion is a negative ion.

Electron Affinity values are usuallynegative since most atoms releaseenergy when they accept an electron.


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6.2.5 Trends in Electron Affinities

Metals on the left of the P.T. tend to

have low Electron Affinities sincethey are usually want to lose anelectron rather than gain one.

Non-Metals on the right of the P.T. tendto have high Electron Affinities sincethey usually want to gain an electronto complete their valence energylevel.

ren n


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. . - ren nElectronegativity

Electronegativity is the tendency of anatom to gain or attract electrons fromanother atom

The +ve protons in the nucleus of oneatom attract the ve electrons from

another atom As the # of protons in the nucleus , theelectronegativity

Therefore, electronegativity across theperiod

The most electronegative element isFluorine

Electronegativity decreases down a groupbecause the electrons are further fromthe nucleus (larger atomic radius) so thenuclear charge (positive charge of

ren n


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. . - ren nElectronegativity

Electronegativity is the ability of anatom to gain electrons (attract or addelectrons)

S f T d


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Summary of Trends1. Atomic Radius: Largest toward bottom left

corner of P.T.

1 .

S f T d


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Summary of TrendsIonization Energy greatest towards Top Right ofIonization Energy greatest towards Top Right of

P.T.P.T.

1 .

S f T d


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Summary of Trends1. Atomic Radius: Largest toward bottom right

corner of P.T.

1 .

2. Ionizat ion Ener gy: Grea t est towa rd top left of P. T.3. Elect ron Affinity: Great est tow ard t op Right of P.

T.


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College of the North Atlantic Qatar

Table Periodic TablePeriodic Table: Map of the Building block of

matter TypeType: Metal, Metalloid, Nonmetal and Noble

Gases

Groupings: Representative or Main, Transition and

Lanthanide/Actanides FamilyFamily: Elements in the same column have

similar chemical properties because ofsimilar valence electrons (outer electrons)

Alkali, Alkaline Earth Metals, Halogens, Noblegases

PeriodPeriod:: Elements in the same row show trendsacross the period due to increasing atomic #(increasing nuclear charge) across the period


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CH 1200 UNIT 7 PART B

7.4.1 VSEPR THEORY V = VALENCE

S = SHELL E = ELECTRON

P = PAIR

R = REPULSION


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7.4.1 ELECTRON PAIRS Atomic Orbitalsmay be filled with 2

electrons. Whenever two electronsoccupy an orbital they are called anelectron pair.

Examples: He (1s2 ) has 2 e- in the 1sorbital. These two electrons form an

electron pair.


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7.4.1 TYPES OF ELECTRONPAIRS Bonding Pair: A pair of electrons ( one

from each atom) used to create abond between the two atoms.

Example: H :H The two electronsbonding the Hydrogen atomstogether in a H2 molecule are a

bonding pair.


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7.4.1 TYPES OF ELECTRONPAIRS Non-Bonding Pair (Lone): A pair of

electrons occupying the same atomicorbital which are not involved inbonding one atom to another.

Example: N 1s2 2s2 2p3 has 2 electronsin the 2s orbital which are not

involved with bonding the Nitrogenatom to other atoms.


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7.4.1 TYPES OF ELECTRONPAIRS Non Bonding (Lone) Pair in Ammonia

NH3

H : N : H H - N - H

l

H H

ypes o ec ron


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. . ypes o ec ronPairs


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7.4.1 VSEPR THEORY

VSEPR Theory suggests that thestructure around a given atom in amolecule is determined by a tendencyto minimize electron-pair repulsions.

All electrons are negative and repeleach other. Pairs of electrons also

repel each other.


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7.4.2 MOLECULAR SHAPES The 3 dimensional shape of a

molecule in space is determinedby the repulsions betweenelectron pairs.

Many molecules have a central

atom around which the otheratoms are arranged in space.

The central atom is usually theatom which can make thehighest number of bonds.


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7.4.2 CENTRAL ATOM

Carbon is thecentral atomin themoleculemethane CH4

Carbon is greyand makes 4bonds.

Hydrogen isreen and


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7.4.2 CENTRAL ATOM

Beryllium is thecentral atom inthe Beryllium

Bromidemolecule.

Beryllium is Grey

and makes 2bonds.

Bromine is Red

and makes 1bond


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7.4.2 CENTRAL ATOM

Boron is thecentral atomin boron

trifluorideBF3

Boron is orange and makes 3

bonds

Fluorine is


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7.4.2 MOLECULAR SHAPES Molecular Shapes are predictable if we

know how manybonds the centralatom can possibly make.

The central atom has 360o of 3D spacearound it. The non-central atoms willspread out as much as possible in this

3D space.


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7.4.2 MOLECULAR SHAPES NOTE The shapes to be described on

the following slides apply to centralatoms around which there are onlyBonding Pairs.

The expected shape will changeslightlyif there are one or more non-bonding

(lone) pairs around the central atom.


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7.4.2 ONE BOND = LINEAR

When two atomsshare onebonding pair of

electrons, theonly possiblemolecular shapeis linear

Example: H2

H : H H - HSeparation 180 o


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7.4.2 TWO BONDS = LINEAR

When the centralatom makestwobonds the

best possiblemolecular shapeis linear.

The bonds are180o

apart.

Example: BeH2 H : Be : H

H Be - H


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7.4.2 THREE BONDS = TRIGONAL

PLANAR ( TRIANGULAR)

When the centralatom makesthree bonds the

best possiblemolecular shapeis trigonalplanar(triangular)

Example: AlCl3 The angle

between the


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7.4.2 Four Bonds = Tetrahedral

When the centralatom makesfourbonds, the

best possibleshape istetrahedral.

Example: CH4

The anglebetween theHydrogen atomsis 109.5o

7 4 3 Structural


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7.4.3 StructuralGeometries The presence of one or more non-

bonding (lone) pairs around a centralatom can change the shape of themolecule slightly.

Non-Bonding (Lone) pairs of electronstend to be closer to the nucleus than

bonding pairs.


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7.4.3 STRUCTURAL PAIRGEOMETRIES N has 3 bonding pairs and 1 non-

bonding (lone) pair of electrons.

The NH3 molecule should be tetrahedral

(109 o) but the presence of the non-bonding (lone) pair reduces this to107 o.

NH3 has a trigonal pyramid shape.

. . ruc ura a r


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. . ruc ura a rGeometries


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7.4.3 STRUCTURAL PAIRGEOMETRIES

O has 2 bondingpairs and 2 non-bonding (lone)pair of

electrons.

The H2O moleculeshould betetrahedral(109.5 o) butthe presence ofthe 2 non-bonding (lone)pair reduces

this to 104.5o

.


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7.4.4 STRUCTURAL VS MOLECULARGEOMETRIES Text Reference

Table 7.4 Page 268 269

This table gives the expected moleculargeometries if there are no non-bonding(lone) pairs.

Also, the changes expected in shape with1, 2, or 3 non-bonding (lone) pairs.

7 4 5 BONDING


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7.4.5 BONDINGCONCEPTS Text Reference

Chapter 10 in McMurry & Fay

10.1 Polar Covalent Bonds

10.2 Intermolecular Forces Please refer to pp 381 390 for additional

information on Bonding Concepts and

Intermolecular Forces.


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7.4.5 MOLECULAR POLARITY Non-Polar Molecules: Molecules in

which the distribution of electrondensity is the same all around themolecule.

Non-Polar molecules result whenidentical atoms or atoms with similar

electronegativity are involved.


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7.4.5 MOLECULAR POLARITY Polar Molecules: Molecules in which

the distribution of electron density isNOT the same all around themolecule.

Polar molecules result when atoms withvery different electronegativities are

involved.


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7.4.5 MOLECULAR POLARITY Polar Molecules can also result from a

distortion of the molecular shape dueto the presence of non-bonding (lone)pairs around the central atom.


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7.4.5 POLARITY OF MOLECULES Bond Dipole: A bond dipole is an arrow

used in a structural diagram to showthe uneven distribution of electrondensity between two bondedatoms.


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7.4.5 POLARITY OF MOLECULES Molecular Dipole: A molecular dipole is

an arrow used in a structural diagramto show the uneven distribution ofelectron density across the entiremolecule.


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7.4.5 POLARITY OF MOLECULES

MolecularDipole ForWater


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7.4.5 POLARITY OF MOLECULES Bond Dipoles vs Molecular Dipoles

A molecule may have more than onedipole. If they are equal and oppositethey cancel and the molecule will NOThave a molecular dipole.

Examples: CCl4

CO2 O = C = O


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7.4.5 POLARITY OF MOLECULES Bond Dipoles vs Molecular Dipoles A molecule may have more than one

dipole. If they are unequal and notopposite they will NOT cancel and themolecule will have a molecular dipole.

Examples: NH3 Trigonal Pyramid

H2O V- Shaped / Bent

7 5 1 INTRAMOLECULAR


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7.5.1 INTRAMOLECULARFORCES IntramolecularForces: Those forces

within an individual molecule whichholds the atoms of the moleculetogether.

The forces of attraction produced bythe sharing of electrons in Bonding

Pairs.

7 5 1 INTERMOLECULAR


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7.5.1 INTERMOLECULARFORCES IntermolecularForces: Those forces of

attraction between molecules in asubstance which holds the moleculestogether as a solid, liquid or gas.

Intermolecular forces are strongest insolids, weaker in liquids and almost

non-existent in gases.

. .FORCES


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FORCES


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7.5.1 VAN de WAALS FORCES Van der Waals Forces: The term

describing all types of intermolecularforces.

London Dispersion Forces

Dipole Dipole Forces

Hydrogen Bonding

Ion Dipole Forces


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7.5.2 DIPOLE-DIPOLE FORCE


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7 5 4 INSTANTANEOUSLY


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7.5.4 INSTANTANEOUSLYINDUCED DIPOLE INTERACTION Instantaneously Induced Dipoles

In some larger symmetrical moleculeswith many electrons the distribution ofcharge may be unsymmetrical at anygiven instant. This results in theformation of a temporary dipolewithin

the molecule. This can affect nearbymolecules.

. .INDUCED DIPOLE INTERACTION


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INDUCED DIPOLE INTERACTION

7 5 4 LONDON DISPERSION


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7.5.4 LONDON DISPERSIONFORCES

London Dispersion Forces:Attractive forces which existbetween all molecules due tothe presence of electrons.

LDF are very weak attractionscompared to ionic or covalentbonds.

LDF are directly related to thenumber of electrons present inthe molecules.

More compact molecules havelower LDF than larger, spreadout molecules.


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7.5.5 ION - ION FORCE

Ion Ion Force:The type ofattraction

which existswithin ioniccrystalswhereoppositelycharged ionsattract.


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7.5.6 ION DIPOLE FORCE Ion Dipole Force: A force of attraction

between a charged ion and the

partial charges on a polar molecule.

Example: When NaCl dissolves in waterthe positive and negative ions formedare attracted to the polar water

molecules.


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7.5.7 HYDROGEN BONDING Hydrogen Bonding: A type of

bonding which arises wheneverHydrogen bonds with a highly-electronegative atom. ( Fluorine/ Oxygen / Nitrogen)

The H F , H O and H Nbonds

tend to be very polar bonds. TheH

electron spends most of its timeclose to the electronegative atom.

7.5.7. HYDROGEN BONDING


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7.5.8 METALLIC BONDING Metallic Bonding: A type of bonding

which exists in solid metals in which

electrons are free to move along thesurface of the metal.

Metallic crystals are 3D arrays of metalcations immersed in a sea ofdelocalized electrons.


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7.5.8 METALLIC BONDING Metallic Bonding accounts for

many of the properties of metals.

Examples: High conductivity of heat and

electricity Ductility & Malleability

7 5 8 NETWORK COVALENT


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7.5.8 NETWORK COVALENTBONDING Network Covalent Bonds: A special

type of covalent bond which results in

very high melting point solids whichare extremely hard.

Examples: Diamond

Silicon Dioxide SiO2


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8.1.1 DEFINE AQUEOUS SOLUTIONSAqueous Solution: Any solution in which

the solute is dissolved in the solvent

water.

8 1 2 DEFINE NON-AQUEOUS


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8.1.2 DEFINE NON AQUEOUSSOLUTIONS Non-Aqueous Solutions: Solutions in

which the solute is dissolved in a

solvent OTHERTHAN water.

8.1.3 LIST TYPES OF SOLUTIONS BY


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8.1.3 LIST TYPES OF SOLUTIONS BYPHYSICAL STATES Gaseous Solutions: Mixtures of gases.

Example: The atmosphere is a mixtureof Nitrogen, Oxygen , Carbon Dioxideand other trace gases.



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8.1.3 LIST TYPES OF SOLUTIONS BYPHYSICAL STATES Liquid Solutions: Mixtures of liquids.

Example: Gasoline is usually a mixtureof different hydrocarbon liquids



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8.1.3 LIST TYPES OF SOLUTIONS BYPHYSICAL STATES Solid Solutions: Mixtures of solids.

Example: Metal alloys are mixtures ofsolid metals. Individual metals aremelted and then poured into onemold to produce an alloy.



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8.1.3 LIST TYPES OF SOLUTIONS BYPHYSICAL STATES

Examples of Alloys:

Brass is a mixture of Cu and Zn

Bronze is a mixture of Cu & Sn

Monel is a mixture of Cu & Ni

8.2.1 DESCRIBE HYDRATION


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8.2.1 DESCRIBE HYDRATIONPROCESS Hydration: The process by which a

solute (solid. liquid or gas) is

dissolved by water.

8 3 1 DEFINE ARRHENIUS ACID


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8.3.1 DEFINE ARRHENIUS ACID Arrhenius Acid: Any substance which

dissolves in water and dissociates to

produce H+(aq) ions.Examples: HCl (g) H

+(aq) + Cl(aq)

HNO3(aq) H +(aq) + NO3 (aq)

8.3.2 DEFINE BRONSTED-LOWRY


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8.3.2 DEFINE BRONSTED LOWRYACID Bronsted-Lowry Acid: Any substance

which can transfer aproton (H+) to

another substance.

Example:

HNO2(aq) + H2O (l) H

3O+

(aq)+ NO

2

-(aq)

8 3 3 DEFINE ARRHENIUS BASE


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8.3.3 DEFINE ARRHENIUS BASE Arrhenius Base: Any substance which

dissolves in water and dissociates to

produce OH-(aq) ions.Examples: NaOH (S) Na

+(aq) + OH(aq)

KOH (s) K +(aq) + OH (aq)

8.3.4 DEFINE BRONSTED-LOWRY


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8.3.4 DEFINE BRONSTED LOWRYBASE Bronsted-Lowry Base: Any substance

which can accept aproton (H+) from

another substance.

Example:

HS

-

(aq) + HF (aq) F-

(aq)+ H

2S

(aq)

8 3 5 PROPERTIES OF ACIDS


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8.3.5 PROPERTIES OF ACIDSAcids

- Solids, Liquids or gases as puresubstances at room temp.

- Soluble in water

- Taste sour

- Form conducting solutions

- Turn Blue Litmus to Red

-

8 3 5 PROPERTIES OF BASES


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8.3.5 PROPERTIES OF BASESBases

- Solids at room temperature

- Soluble in water

- Taste bitter

- Form conducting solutions

- Turn Red Litmus to Blue.

8 3 6 PROCESS OF IONIZATION


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8.3.6 PROCESS OF IONIZATIONIonization: The process by which neutral

atoms lose or gain one or more

electrons to form ions.

Examples:

Na Na + + 1 e

F + 1 e - F -

8 3 6 PROCESS OF DISSOCIATION


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8.3.6 PROCESS OF DISSOCIATIONDissociation: The process by which a

molecule dissociates into positive and

negative ions.

Examples: CaCl2(s) Ca

2+ (aq) + 2 Cl(aq)

Ba(OH)2(s) Ba2+ (aq) + 2 OH

(aq)

8 3 7 DEFINE ELECTYROLYTE


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8.3.7 DEFINE ELECTYROLYTEElectrolyte: Any substance which

dissolves in water and produces a

solution which conductselectricity.Electrolytes provide mobile aqueousions in solution which conduct current.

8 3 7 DEFINE NON ELECTROLYTE


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8.3.7 DEFINE NON-ELECTROLYTENon-Electrolyte: Any substance which

dissolves in water and forms a solution

which does NOT conduct electricity.Molecular compounds are typical non-

electrolytes since they do not form ionsin solution.

8 3 8 DEFINE AMPHIPROTIC


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8.3.8 DEFINE AMPHIPROTICAmphoteric: A term which describes any

substance which can behave like a B-L

Acid or a B-L base.

Molecules or Anions with availableHydrogen may be amphoteric.


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8.3.8 Amphoteric ExamplesHydrogen Carbonate Ion HCO3

-

Acting as a B-L Acid:

HCO3-(aq) + OH

-(aq) CO 3

2-(aq) +

H2O (l)

Acting as a B-L Base:HCO3

-(aq) + H3O+

(aq) H 2CO3(aq) +

H2O(l)

8 3 9 DEFINE POLYPROTIC ACID


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8.3.9 DEFINE POLYPROTIC ACIDPolyprotic Acid: Any acid which has more

than one Hydrogen atom that

dissociates.Examples:

H2SO4(aq) yields 2 H+ ions

H3PO

4(aq)

yields 3 H+ ions

8.3.10 DEFINE NEUTRALIZATION


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REACTIONNeutralization Reaction: A reaction

between an acid and a base which

produces water and a salt.Example:

HCl (aq) + NaOH (aq) HOH (l)+

NaCl (aq)

8.3.11 FORMULAS OF COMMON


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ACIDSCommon Mineral Acids Hydrochloric acid HCl (aq) Nitric Acid HNO3(aq) Sulfuric Acid H2SO4(aq)

8.3.11 FORMULAS OF COMMON


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BASESCommon Bases: Sodium Hydroxide NaOH Lithium Hydroxide LiOH Potassium Hydroxide KOH

Barium Hydroxide Ba(OH)2

Strontium Hydroxide Sr(OH)2

Documents

Basic Chemistry 04