1. what is an atom made of (think of what you have learned in other classes)? day 3 10-16

1. What is an atom made of (think of what you have learned in other classes)?

Day 3 10-16

Atoms, Molecules, and Ions

Chapter 2

Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

1. Brainstorm with your discussion partner: everything you know about an atom.

2. Look up the definition in the back of the book, and record it

3. Draw an atom of hydrogen. Draw an atom of oxygen. (There is no wrong answer at this time… Draw something!)

ATOM Neutron

Proton

Electron

(0)

(+)

(-)

Nucleus

When and Why do we use models?

In chemistry:

I. Early Greeks• Everything is made

of 4 elements: _______________, _______________,

_______________, and _______________.

• These combine and interact to make everything.

EarthEarthFireFireAirAir

WaterWater

Continuous matter – accepted for nearly 200 years

-Aristotle & Plato

400 B.C. Basic particle = an atom – “indivisible”

-Democritus

The Atom in Ancient Greece

Sparks of Knowledge - 1700’s

• Law of conservation of matter states that during any ________ or _________ process, matter is neither _______ nor ________

• In a reaction, the mass of the reactants ___________ the mass of the products.

chemicalchemical physicalphysicalcreatedcreated destroyeddestroyed

E Q U A L SE Q U A L S

Sparks of Knowledge - 1700’s Joseph Proust

• Developed the law of definite proportions or constant composition which

states that the _____ ratio of elements in a compound is always _________.

• Examples:

massmassthe samethe same

Water 1g H : 8g Water 1g H : 8g OOCarbon Dioxide 3g C : Carbon Dioxide 3g C : 8g O8g O

Law of Conservation of Mass – Mass is neither destroyed nor created

Reactantsmass = Productsmass

Law of Definite Proportions – A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample

Lasting Laws

Dalton’s Atomic Theory (1800s):

1. All matter is composed of extremely small particles called atoms.

2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.

3. Atoms cannot be subdivided, created, or destroyed.

4. Atoms of different elements combine in simple whole-numbered ratios to form chemical compounds.

5. In chemical reactions, atoms are combined, separated, or rearranged.

Modern Atomic Theory:

1. All matter is composed of extremely small particles called atoms.

2. A GIVEN ELEMENT CAN HAVE ATOMS WITH DIFFERENT MASSES

3. ATOMS ARE DIVISIBLE INTO EVEN SMALLER PARTICLES.

4. Atoms of different elements combine in simple whole-numbered ratios to form chemical compounds.

5. In chemical reactions, atoms are combined, separated, or rearranged.

Dalton’s Atomic Theory (1800s):

1.All matter is composed of extremely small particles called atoms.

Modern Atomic Theory:

1.All matter is composed of extremely small particles called atoms.

Dalton’s Atomic Theory (1800s):

2.Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.

Modern Atomic Theory:

2.A GIVEN ELEMENT CAN HAVE ATOMS WITH DIFFERENT MASSES

Dalton’s Atomic Theory (1800s):

3.Atoms cannot be subdivided, created, or destroyed.

Modern Atomic Theory:

3.ATOMS ARE DIVISIBLE INTO EVEN SMALLER PARTICLES.

Dalton’s Atomic Theory (1800s):

4.Atoms of different elements combine in simple whole-numbered ratios to form chemical compounds.

Modern Atomic Theory:

4.Atoms of different elements combine in simple whole-numbered ratios to form chemical coms.

Dalton’s Atomic Theory (1800s):

5. In chemical reactions, atoms are combined, separated, or rearranged.

Modern Atomic Theory:

5. In chemical reactions, atoms are combined, separated, or rearranged.

1. One thing Dalton had right was …

2. One thing Dalton was wrong about was…

Pd 3 Day 4 10-19

We have Atoms! – 1803-09 John Dalton

Atom’s Appearance• __________• __________• _____ ________• Different __________ • Different __________• Different __________

SolidSolidSphereSphereLike marblesLike marbles

sizessizescolorscolors

weightsweights

19

Dalton’s Atomic Theory (1808)1. Elements are composed of extremely small particles

called atoms.

2. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements.

3. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction.

4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction.

Review sections 2.1 and 2.2 and complete #s 2.1-2.8 on page 68.

- due Tuesday

20

1. One thing Dalton had right was …

2. One thing Dalton was wrong about was…

Pd 1 Day 4 10-19

1. One thing Dalton had right was …

2. One thing Dalton was wrong about was…

Pd 8 Day 5 10-20

23

8 X2Y16 X 8 Y+

Law of Conservation of Mass

We have Atoms! - 1850’s • Cathode Ray Tubes

–Invented and studied in the mid-1800’s

–A beam of cathode ray particles . . .

• applies ________,

• makes the temperature of objects _________,

• and is affected by ___________ and ___________ fields.

forceforce

warmerwarmerelectricelectric

magneticmagnetic

We have Atoms! - 1850’s Cathode Ray Tubes

We have Atoms! - 1850’s Cathode Ray Tubes

We have Atoms! – 1897 J. J. Thomson

• Studied _________ ______.

• All metals create _______________________.

• Cathode ray particles are _____ and _________ _________.

• Thomson coined the _______ as another name for ______ ___ _________.

cathode cathode raysrays

identical cathode identical cathode raysrays tinytiny

negatively negatively chargedcharged

electronelectron

particlesparticlescathode cathode ray ray

We have Atoms! – 1897 J. J. Thomson

solidsolidspheresphere

negative electronsnegative electronspositive positive surroundingssurroundings

• Atom’s appearance:

–______

–_______–Parts

•________________

•_______ ___________

29

Thomson’s Model

30

J.J. Thomson, measured mass/charge of e-

(1906 Nobel Prize in Physics)

Cathode Ray Tube

31

Cathode Ray Tube

Review sections 2.1 and 2.2 and complete #s 2.1-2.8 on page 68.

- due Tuesday

32

33

e- charge = -1.60 x 10-19 C

Thomson’s charge/mass of e- = -1.76 x 108 C/g

e- mass = 9.10 x 10-28 g

Measured mass of e-

(1923 Nobel Prize in Physics)

Millikan’s Experiment

Review sections 2.1 and 2.2 and complete #s 2.1-2.8 on page 68.

- due Tuesday

34

• Investigated _________ and its behaviors.

• The chemical properties of radioactive elements _______ as radiation is emitted.

radiationradiation

changechange

We have Atoms! – 1898Marie Skłodowska Curie and Pierre Curie

• Isolated and discovered the elements _________ and ________.

poloniumpoloniumradiumradium

We have Atoms! – 1898Marie Skłodowska Curie and Pierre Curie

We have Atoms! – 1899 Ernest Rutherford

• Analyzed ________ to determine its ___________.

–______________, , are fairly heavy with a ________ charge.

–_____________, , are very tiny with a ________ charge.

radiationradiationcompositioncompositionAlpha particlesAlpha particles

positivepositiveBeta particlesBeta particles

negativenegative

We have Atoms! – 1899 Ernest Rutherford

–_____________, , are very tiny with a ________ charge.

–____________, , have no mass with a _______ charge.

Beta particlesBeta particlesnegativenegative

Gamma raysGamma raysneutralneutral

Radiation Penetration

Radiation Separation

• What does this separation pattern indicate about charges and masses of these types of radiation?–α: ______ mass and ________ charge

–β: ________ mass and _________ charge

largerlarger positivepositivesmallersmaller negativenegative

• What does this separation pattern indicate about charges and masses of these types of radiation?–β: ________ mass and _________

charge

–γ: ___ mass and ___ charge

smallersmaller negativenegative

nono nono

Radiation Separation

42

(uranium compound)

Types of Radioactivity

Atoms defy what we thought we knew! 1909-1910

• Ernest Rutherford’s–Experiment applet #1

–Experiment applet #2

• Ernest Rutherford’s

Atom’s appearance:– ____________________– ____________________________________– ________________________________

Hollow Hollow spheresphereDense center with positive chargeDense center with positive charge

Electrons occupy empty spaceElectrons occupy empty space

45

1. atoms positive charge is concentrated in the nucleus2. proton (p) has opposite (+) charge of electron (-)3. mass of p is 1840 x mass of e- (1.67 x 10-24 g)

particle velocity ~ 1.4 x 107 m/s(~5% speed of light)

(1908 Nobel Prize in Chemistry)

Rutherford’s Experiment

46

atomic radius ~ 100 pm = 1 x 10-10 m

nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m

Rutherford’s Model of the Atom

“If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.”

Atoms defy what we thought we knew! 1932

• James Chadwick discovered _________

–Have no _______

–Located in the atom’s

_______

neutronsneutronschargecharge

nucleusnucleus

48

Chadwick’s Experiment (1932)(1935 Noble Prize in Physics)

H atoms: 1 p; He atoms: 2 p

mass He/mass H should = 2

measured mass He/mass H = 4

+ 9Be 1n + 12C + energy

neutron (n) is neutral (charge = 0)

n mass ~ p mass = 1.67 x 10-24 g

Parts of Atoms, Isotopes, and Ions

Fundamental parts of atoms

COMPONENTCHARGE

APPROXIMATE MASS (in amu)

LOCATION IN ATOM

electronelectron - 1- 1 0.000550.00055

shellshellss

protonproton + 1+ 1 1.0 1.0 nucleusnucleus

neutronneutron 00 1.000551.00055 nucleusnucleus

Parts of Atoms, Isotopes, + Ions

–Individual _____ of any element

–Represented with a _______ symbol

XXAA

ZZ

CC

atomatomnuclearnuclear

element symbolelement symbolmass numbermass number

atomic numberatomic number

•X = _______ _______

•A = ____ _______

•Z = ______ _______

Parts of Atoms, Isotopes, and Ions

•X = _______ _______

•A = ____ _______

•Z = ______ _______

= ______ _______

•C = ______ __ __ ___

•# neutrons = ___ - ___

•# electrons = ___ - ___

XXAA

ZZ

CCelement symbolelement symbolmass numbermass number

atomic numberatomic numberprotons numberprotons number

charge as an ioncharge as an ionAA ZZ

ZZ CC

Parts of Atoms, Isotopes, and Ions

XX1919 1-1-Mass number?Mass number?

Atomic number?Atomic number?

# of neutrons?# of neutrons?

Element?Element?

Charge?Charge?

# of electrons?# of electrons?

99

Parts of Atoms, Isotopes, and Ions

• Isotopes

–Atoms of one element

•Same number of _______

•Different numbers of ________

•Different ______

protonsprotonsneutronsneutrons

massesmasses

Parts of Atoms, Isotopes, and Ions

• Ions

–________ atoms that have gained or lost ________

–Examples

• Iron loses 2 electrons

•Chlorine gains 1 electron

ChargedChargedelectronselectrons

FeFe +2 +2

ClCl -1 -1

55

mass p ≈ mass n ≈ 1840 x mass e-

56

Atomic number (Z) = number of protons in nucleus

Mass number (A) = number of protons + number of neutrons

= atomic number (Z) + number of neutrons

Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei

XAZ

H11 H (D)2

1 H (T)31

U23592 U238

92

Mass Number

Atomic NumberElement Symbol

Atomic Number, Mass Number, and Isotopes

57

The Isotopes of Hydrogen

Example 2.1

Give the number of protons, neutrons, and electrons in each of the following species:

(a)

(b)

(c)

(d) carbon-14

Example 2.1

Solution

(a) The atomic number is 11, so there are 11 protons. The mass number is 20, so the number of neutrons is 20 − 11 = 9. The number of electrons is the same as the number of protons; that is, 11.

(b) The atomic number is the same as that in (a), or 11. The mass number is 22, so the number of neutrons is 22 − 11 = 11. The number of electrons is 11. Note that the species in (a) and (b) are chemically similar isotopes of sodium.

Example 2.1

(c) The atomic number of O (oxygen) is 8, so there are 8 protons. The mass number is 17, so there are 17 − 8 = 9 neutrons. There are 8 electrons.

(d) Carbon-14 can also be represented as 14C. The atomic number of carbon is 6, so there are 14 − 6 = 8 neutrons. The number of electrons is 6.

Example

Atomic number = # of protons = # of electrons (neutral atom)

Atomic mass

Carbon has an atomic number of 6, which means it has 6 ________ in its nucleus and 6 __________ orbiting around the nucleus.

protonselectrons

Organization of Your Periodic Table

He, Ne, and Ar are in the same ___________.

C, N, and O are in the same ___________.

P E R I O D SGROUPS

metalloids

Metals

Excellent conductors (Sea of electrons) even as solids

Mobile electrons

Metals are malleable and ductile

Can be pounded into sheet

Can be pulled into wire

Alkali Metals

Similar properties:

Group # 1

Reactive

Good Conductors

Alkaline Earth Metals

Similar properties:

Group # 2

Reactive – but less than alkali

Good Conductors

Noble Gases

Similar properties:

Group # 18

NON-reactive

Gases

Full outer shells!!!

Halogens

Similar properties:

Group # 17

Very reactive

Fluorine and Chlorine = gases

7 valence electrons

Transition metalsSimilar properties:

Generally least reactive metals

Tricky valence (can change from the expected)

Mostly form +1, +2, and +3 ions

71

The Modern Periodic Table

PeriodG

roup

Alkali M

etal

Noble G

as

Halogen

Alkali E

arth Metal

72

Chemistry In ActionNatural abundance of elements in Earth’s crust:

Natural abundance of elements in human body:

73

A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical forces.

H2 H2O NH3 CH4

A diatomic molecule contains only two atoms:

H2, N2, O2, Br2, HCl, CO

A polyatomic molecule contains more than two atoms:

O3, H2O, NH3, CH4

diatomic elements

74

An ion is an atom, or group of atoms, that has a net positive or negative charge.

cation – ion with a positive chargeIf a neutral atom loses one or more electronsit becomes a cation.

anion – ion with a negative chargeIf a neutral atom gains one or more electronsit becomes an anion.

Na 11 protons11 electrons Na+ 11 protons

10 electrons

Cl 17 protons17 electrons Cl-

17 protons18 electrons

75

A monatomic ion contains only one atom:

A polyatomic ion contains more than one atom:

Na+, Cl-, Ca2+, O2-, Al3+, N3-

OH-, CN-, NH4+, NO3

-

76

Common Ions Shown on the Periodic Table

77

Formulas and Models

78

A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance.

An empirical formula shows the simplest whole-number ratio of the atoms in a substance.

H2OH2O

molecular empirical

C6H12O6 CH2O

O3 O

N2H4 NH2

Example 2.2

Write the molecular formula of methanol, an organic solvent and antifreeze, from its ball-and-stick model, shown below.

Example 2.2

Solution

Refer to the labels (also see back endpapers).

There are four H atoms, one C atom, and one O atom. Therefore, the molecular formula is CH4O.

However, the standard way of writing the molecular formula for methanol is CH3OH because it shows how the atoms are joined

in the molecule.

Example 2.3

Write the empirical formulas for the following molecules:

(a)acetylene (C2H2), which is used in welding torches

(b)glucose (C6H12O6), a substance known as blood sugar

(c)nitrous oxide (N2O), a gas that is used as an anesthetic gas (“laughing gas”) and as an aerosol propellant for whipped creams.

Example 2.3

Strategy

Recall that to write the empirical formula, the subscripts in the molecular formula must be converted to the smallest possible whole numbers.

Example 2.3

Solution

(a)There are two carbon atoms and two hydrogen atoms in acetylene. Dividing the subscripts by 2, we obtain the empirical formula CH.

(b) In glucose there are 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms. Dividing the subscripts by 6, we obtain the empirical formula CH2O. Note that if we had divided the subscripts by 3, we would have obtained the formula C2H4O2.

Although the ratio of carbon to hydrogen to oxygen atoms in C2H4O2 is the same as that in C6H12O6 (1:2:1), C2H4O2 is not the

simplest formula because its subscripts are not in the smallest whole-number ratio.

Example 2.3

(c) Because the subscripts in N2O are already the smallest

possible whole numbers, the empirical formula for nitrous oxide is the same as its molecular formula.

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Ionic compounds consist of a combination of cations and anions.

• The formula is usually the same as the empirical formula.

• The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero.

The ionic compound NaCl

86

The most reactive metals (green) and the most reactive nonmetals (blue) combine to form ionic compounds.

87

Formulas of Ionic Compounds

Al2O3

2 x +3 = +6 3 x -2 = -6

Al3+ O2-

CaBr2

1 x +2 = +2 2 x -1 = -2

Ca2+ Br-

Na2CO3

2 x +1 = +2 1 x -2 = -2

Na+ CO32-

Example 2.4

Write the formula of magnesium nitride, containing the Mg2+ and N3− ions.

When magnesium burns in air, it forms both magnesium oxide and magnesium nitride.

Example 2.4

Strategy Our guide for writing formulas for ionic compounds is electrical neutrality; that is, the total charge on the cation(s) must be equal to the total charge on the anion(s).

Because the charges on the Mg2+ and N3− ions are not equal, we know the formula cannot be MgN.

Instead, we write the formula as MgxNy, where x and y are

subscripts to be determined.

Example 2.4

Solution To satisfy electrical neutrality, the following relationship must hold:

(+2)x + (−3)y = 0

Solving, we obtain x/y = 3/2. Setting x = 3 and y = 2, we write

Check The subscripts are reduced to the smallest whole- number ratio of the atoms because the chemical formula of an ionic compound is usually its empirical formula.

91

Chemical Nomenclature

• Ionic Compounds– Often a metal + nonmetal– Anion (nonmetal), add “-ide” to element name

BaCl2 barium chloride

K2O potassium oxide

Mg(OH)2 magnesium hydroxide

KNO3 potassium nitrate

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• Transition metal ionic compounds– indicate charge on metal with Roman numerals

FeCl2 2 Cl- -2 so Fe is +2 iron(II) chloride

FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride

Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide

93

94

Example 2.5

Name the following compounds:

(a)Cu(NO3)2

(b)KH2PO4

(c)NH4ClO3

Example 2.5

Strategy Note that the compounds in (a) and (b) contain both metal and nonmetal atoms, so we expect them to be ionic compounds.

There are no metal atoms in (c) but there is an ammonium group, which bears a positive charge. So NH4ClO3 is also an

ionic compound.

Our reference for the names of cations and anions is Table 2.3.

Keep in mind that if a metal atom can form cations of different charges (see Figure 2.11), we need to use the Stock system.

Example 2.5

Solution

(a)The nitrate ion ( ) bears one negative charge, so the copper ion must have two positive charges. Because copper forms both Cu+ and Cu2+ ions, we need to use the Stock system and call the compound copper(II) nitrate.

(b)The cation is K+ and the anion is (dihydrogen phosphate). Because potassium only forms one type of ion (K+), there is no need to use potassium(I) in the name. The compound is potassium dihydrogen phosphate.

(c) The cation is (ammonium ion) and the anion is . The compound is ammonium chlorate.

Example 2.6

Write chemical formulas for the following compounds:

(a)mercury(I) nitrite

(b)cesium sulfide

(c)calcium phosphate

Example 2.6

Strategy

We refer to Table 2.3 for the formulas of cations and anions.

Recall that the Roman numerals in the Stock system provide useful information about the charges of the cation.

Example 2.6

Solution

(a)The Roman numeral shows that the mercury ion bears a +1 charge. According to Table 2.3, however, the mercury(I) ion is diatomic (that is, ) and the nitrite ion is . Therefore, the formula is Hg2(NO2)2.

(b)Each sulfide ion bears two negative charges, and each cesium ion bears one positive charge (cesium is in Group 1A, as is sodium). Therefore, the formula is Cs2S.

Example 2.6

(c) Each calcium ion (Ca2+) bears two positive charges, and each phosphate ion ( ) bears three negative charges.

To make the sum of the charges equal zero, we must adjust the numbers of cations and anions:

3(+2) + 2(−3) = 0

Thus, the formula is Ca3(PO4)2.

102

• Molecular compounds− Nonmetals or nonmetals + metalloids− Common names

− H2O, NH3, CH4

− Element furthest to the left in a period and closest to the bottom of a group on periodic table is placed first in formula

− If more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom

− Last element name ends in -ide

103

HI hydrogen iodide

NF3 nitrogen trifluoride

SO2 sulfur dioxide

N2Cl4 dinitrogen tetrachloride

NO2 nitrogen dioxide

N2O dinitrogen monoxide

Molecular Compounds

Example 2.7

Name the following molecular compounds:

(a)SiCl4

(b)P4O10

Example 2.7

Strategy

We refer to Table 2.4 for prefixes.

In (a) there is only one Si atom so we do not use the prefix “mono.”

Solution

(a)Because there are four chlorine atoms present, the compound is silicon tetrachloride.

(b)There are four phosphorus atoms and ten oxygen atoms present, so the compound is tetraphosphorus decoxide. Note that the “a” is omitted in “deca.”

Example 2.8

Write chemical formulas for the following molecular compounds:

(a)carbon disulfide

(b) disilicon hexabromide

Example 2.8

Strategy

Here we need to convert prefixes to numbers of atoms (see Table 2.4).

Because there is no prefix for carbon in (a), it means that there is only one carbon atom present.

Solution

(a)Because there are two sulfur atoms and one carbon atom present, the formula is CS2.

(b) There are two silicon atoms and six bromine atoms present, so the formula is Si2Br6.

108

109

An acid can be defined as a substance that yields hydrogen ions (H+) when dissolved in water.

For example: HCl gas and HCl in water

•Pure substance, hydrogen chloride

•Dissolved in water (H3O+ and Cl−), hydrochloric acid

110

111

An oxoacid is an acid that contains hydrogen, oxygen, and another element.

HNO3 nitric acid

H2CO3 carbonic acid

H3PO4 phosphoric acid

112

Naming Oxoacids and Oxoanions

113

The rules for naming oxoanions, anions of oxoacids, are as follows:1. When all the H ions are removed from the

“-ic” acid, the anion’s name ends with “-ate.” 2. When all the H ions are removed from the

“-ous” acid, the anion’s name ends with “-ite.”3. The names of anions in which one or more

but not all the hydrogen ions have been removed must indicate the number of H ions present.

For example:– H2PO4

- dihydrogen phosphate– HPO4 2- hydrogen phosphate– PO4

3- phosphate

114

Example 2.9

Name the following oxoacid and oxoanion:

(a)H3PO3

(b)

Example 2.9

Strategy To name the acid in (a), we first identify the reference acid, whose name ends with “ic,” as shown in Figure 2.15.

In (b), we need to convert the anion to its parent acid shown in Table 2.6.

Solution

•We start with our reference acid, phosphoric acid (H3PO4). Because H3PO3 has one fewer O atom, it is called phosphorous acid.

•The parent acid is HIO4. Because the acid has one more O atom than our reference iodic acid (HIO3), it is called periodic acid. Therefore, the anion derived from HIO4 is called periodate.

117

A base can be defined as a substance that yields hydroxide ions (OH-) when dissolved in water.

NaOH sodium hydroxide

KOH potassium hydroxide

Ba(OH)2 barium hydroxide

118

Hydrates are compounds that have a specific number of water molecules attached to them.

BaCl2•2H2O

LiCl•H2O

MgSO4•7H2O

Sr(NO3)2 •4H2O

barium chloride dihydrate

lithium chloride monohydrate

magnesium sulfate heptahydrate

strontium nitrate tetrahydrate

CuSO4•5H2O CuSO4

119

120

Organic chemistry is the branch of chemistry that deals with carbon compounds.

C

H

H

H OH C

H

H

H NH2 C

H

H

H C OH

O

methanol methylamine acetic acid

Functional Groups:

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