Ch 2. Atoms and Elements Atom Nucleus (Rutherford’s experiment) Electrons (Thomson’s experiment)

Ch 2. Atoms and Elements

Atom

Nucleus (Rutherford’s experiment)

Electrons (Thomson’s experiment)

Atomic Nucleus

Nucleus carries positive charges

Each electron carries one negative charge

Nucleus carries almost all the mass of an atom

Atom

Nucleus

Electrons (each carries a negative charge)

Protons (each carries a positive charge)

Neutrons (neutral)

number of electrons = number of protons

= atomic number

Atoms are neutral

Atoms that have the same number of protons belong to

one kind of element.

isotopes: same number of protons, different number of neutrons

mass of a proton ≈ mass of a neutron >> mass of an electron

number of protons + number of neutrons = mass number

XA

ZChemical symbol

Mass number

Atomic number

or X-A

Practice – Complete the table

11

Al2713

Tro: Chemistry: A Molecular Approach, 2/e


12

Al2713

C136

Mo9642

Cs13355


No change occurs inside a nucleus in chemistry

Atoms can lose or gain electrons

Na − e− Na+ positive ion = cation

Cl + e− Cl− negative ion = anion

Mg − 2e− Mg2+

O + 2e− O2−


14

3Al



15

2Mg

2S

Br

3Al


SymbolNumber of Protons in

Nucleus

Number of Neutrons in Nucleus

Number of Electrons

Net charge

87Rb+

16 18 2−

36 28 1+

SymbolNumber of Protons in

Nucleus

Number of Neutrons in Nucleus

Number of Electrons

Net charge

87Rb+ 37 50 36 1+

32S2− 16 18 18 2−

65Cu+ 29 36 28 1+

isotopes: same number of protons, different number of neutrons

mass of a proton ≈ mass of a neutron >> mass of an electron

number of protons + number of neutrons = mass number

1H 1.6735 x 10−24 g

16O 2.6560 x 10−23 g

One atomic mass unit (amu) is defined as 1/12 ofthe mass of a 12C atom.

12C atom: 6 protons, 6 neutrons, 6 electrons.

1 amu = 1.6605 x 10−24 g

1H 1.6735 x 10−24 g

16O 2.6560 x 10−23 g

1 amu = 1.6605 x 10−24 g

amu 1.0078g101.6605

amu 1g10 1.6735

2424

2324

1 amu2.6560 10 g 15.995 amu

1.6605 10 g

21

Mass Spectrometer


Mass Spectrum of Natural Copper

69.09 % 63Cu

30.91 % 65Cu

natural abundance:percent of an isotopein nature


69.09 % 63Cu: 62.93 amu

30.91 % 65Cu: 64.93 amu

Average atomic mass of Cu = ?

How to find the average?

1, 1, 1, 1, 2

Average = (1 + 1 + 1 + 1 + 2) / 5 = 1.2

2.1

4.08.0

2%201%80

25

11

5

45

2114

5

21111


69.09 % 63Cu: 62.93 amu

30.91 % 65Cu: 64.93 amu

Average atomic mass of Cu = ?

63.55 amu

listed in the periodic table

35Cl 34.967 amu 75.78 %

37Cl 36.966 amu 24.22 %

Average atomic mass of Cl

= 34.967 amu x 75.78 % + 36.966 amu x 24.22 %

= 35.45 amu

listed in the periodic table

mass of a proton ≈ mass of a neutron ≈ 1 amu

mass number of an isotope ≈ atomic mass of the isotope in amu

35Cl 34.967 amu 75.78 %

37Cl 36.966 amu 24.22 %

One atomic mass unit (amu) is defined as 1/12 ofthe mass of a 12C atom.

The number of carbon atoms in exactly 12 g of 12Cis called Avogadro’s number: 6.022 x 1023

One Avogadro’s number of particles is called a mole.

1 mol = 6.022 x 1023 particles

1 pair = 2 particles 1 dozen = 12 particles

(exact number)

Example 2.6, page 67

Calculate the number of atoms in 2.45 mol of Cu.


Practice 2.6, page 67

A pure Ag ring contains 2.80 x 1022 Ag atoms. How many moles of Ag does it contain?


For an element X, its atomic mass is x amu.

What is the mass of 1 mol of X in grams?

The mass of 1 mol of X is x g.

The molar mass of an element is the mass in grams per mole of the element.

Unit: g/mol

Two ways to find molar mass

1) Read from the periodic table

moles

grams in massmassmolar

2) Use the definition of molar mass:

(recall d = m/V)


moles

grams in massmassmolar

Unit: g/mol

molar mass and Avogadro’s number are exact numbers

A piece of Cu has a mass of 200. g. How many copper

atoms are present?


mass in gramsmolar mass

moles

A silicon chip has a mass of 5.68 mg. How many silicon

atoms are present in the chip?



moles

Compute the mass in grams of a sample of six

Americium atoms.



moles

Calculate the number of moles in a sample of cobalt (Co)

containing 5.00 x 1020 atoms and the mass of the sample.



moles

40

Mendeleev• Ordered elements by atomic mass• Saw a repeating pattern of properties • Periodic Law – when the elements are arranged

in order of increasing atomic mass, certain sets of properties recur periodically

• Put elements with similar properties in the same column

• Used pattern to predict properties of undiscovered elements

• Where atomic mass order did not fit other properties, he re-ordered by other properties– Te & I


41

Periodic Pattern


Periodic Patterns

Li Be B C N O F

Na Mg Al Si P S Cl

K Ca

H

NM H2Oa/b

H21.0

M Li2Ob

LiH6.9

M Na2Ob

NaH23.0

M K2Ob

KH39.1

M BeOa/b

BeH29.0

M MgOb

MgH224.3

M CaOb

CaH240.1

NM B2O3

a

BH310.8

M Al2O3

a/b

AlH327.0

NM CO2

a

CH412.0

M/NM SiO2

a

SiH428.1

NM N2O5

a

NH314.0

NM P4O10

a

PH331.0

NM

H2S32.1

SO3

a

NM

H2O16.0

O2

NM Cl2O7

a

HCl35.5

NM

HF19.0

a = acidic oxide, b = basic oxide, a/b = amphoteric oxide

M = metal, NM = nonmetal, M/NM = metalloid

42Tro: Chemistry: A Molecular Approach, 2/e

43

About ¾ of the elements are classified as metals. They have a reflective surface, conduct heat and electricity better than other elements, and are malleable and ductile

Most of the remaining elements are classified as nonmetals. Their solids have a non-reflective surface, do not conduct heat and electricity well, and are brittle.

A few elements are classified as metalloids. Their solids have some characteristics of metals and some of nonmetals.


44

Metals• Solids at room temperature, except Hg• Reflective surface

– shiny• Conduct heat• Conduct electricity• Malleable

– can be shaped• Ductile

– can be drawn or pulled into wires• Lose electrons and form cations in

reactions• About 75% of the elements are metals• Lower left on the table


45

Nonmetals• Found in all three states

• Poor conductors of heat

• Poor conductors of electricity

• Solids are brittle

• Gain electrons in reactions to become anions

• Upper right on the table– except H

Sulfur, S(s)

Bromine, Br2(l)

Chlorine, Cl2(g)


46

Metalloids

• Show some properties of metals and some of nonmetals

• Also known as semiconductors Properties of Silicon

shinyconducts electricity

does not conduct heat wellbrittle


47

The Modern Periodic Table

• Elements with similar chemical and physical properties are in the same column

• Columns are called Groups or Families– designated by a number and letter at top

• Rows are called Periods• Each period shows the pattern of

properties repeated in the next period


48

The Modern Periodic Table

• Main group = representative elements = “A” groups

• Transition elements = “B” groups– all metals

• Bottom rows = inner transition elements = rare earth elements– metals– really belong in Period 6 & 7


50

= Alkali metals

= Alkali earth metals

= Noble gases

= Halogens

= Lanthanides

= Actinides

= Transition metals


51

Important Groups - Hydrogen

• Nonmetal • Colorless, diatomic gas

– very low melting point and density

• Reacts with nonmetals to form molecular compounds– HCl is acidic gas

– H2O is a liquid

• Reacts with metals to form hydrides– metal hydrides react with water to form H2

• HX dissolves in water to form acids


lithium

sodium

potassium

rubidium

cesium

52

Important Groups – Alkali Metals• Group IA = Alkali Metals• Hydrogen usually placed here,

though it doesn’t really belong• Soft, low melting points, low

density• Flame tests Li = red, Na =

yellow, K = violet• Very reactive, never find

uncombined in nature• Tend to form water-soluble

compounds, therefore salt is crystallized from seawater then molten salt is electrolyzed

• colorless solutions• React with water to form basic

(alkaline) solutions and H2

2 Na + 2 H2O 2 NaOH + H2 • releases a lot of heat


53

Important Groups – Alkali Earth Metals• Group IIA = Alkali earth metals

• Harder, higher melting, and denser than alkali metals – Mg alloys used as structural materials

• Flame tests Ca = red, Sr = red, Ba = green• Reactive, but less than corresponding alkali metal• Form stable, insoluble oxides from which they are

normally extracted• Oxides are basic = alkaline earth

• Reactivity with water to form H2 – Be = none; Mg = steam; Ca, Sr, Ba = cold water


54

Important Groups – Halogens• Group VIIA = halogens• Nonmetals

• F2 and Cl2 gases; Br2 liquid; I2 solid

• All diatomic• Very reactive

• Cl2, Br2 react slowly with water

Br2 + H2O HBr + HOBr

• React with metals to form ionic compounds

• HX all acids– HF weak < HCl < HBr < HI

bromine

iodine

chlorine

fluorine

astatine


55

Important Groups – Noble Gases

• Group VIIIA = Noble Gases• All gases at room temperature

– very low melting and boiling points

• Very unreactive, practically inert• Very hard to remove electron from or give

electron to


56

Ion Charge and the Periodic Table• The charge on an ion can often be

determined from an element’s position on the Periodic Table

• Metals always form positively charged cations

• For many main group metals, the charge = the group number

• Nonmetals form negatively charged anions• For nonmetals, the charge = the group

number − 8


Practice – What is the charge on each of the following ions?

58

• potassium cation

• sulfide anion

• calcium cation

• bromide anion

• aluminum cation

K+

S2−

Ca2+

Br−

Al3+


Law of Conservation of Mass

59

Antoine Lavoisier1743-1794


• In a chemical reaction, matter is neither created nor destroyed

• Total mass of the materials you have before the reaction must equal the total mass of the materials you have at the endtotal mass of reactants =

total mass of products

Reaction of Sodium with Chlorine to Make Sodium

Chloride

60

7.7 g Na + 11.9 g Cl2 19.6 g NaCl


• The mass of sodium and chlorine used is determined by the number of atoms that combine

• Because only whole atoms combine and atoms are not changed or destroyed in the process, the mass of sodium chloride made must equal the total mass of sodium and chlorine atoms that combine together

Law of Definite Proportions

• All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements

61

Joseph Proust1754-1826


62

Proportions in Sodium ChlorideA 100.0 g sample of sodium chloride contains 39.3 g of sodium and 60.7 g of chlorine

A 200.0 g sample of sodium chloride contains 78.6 g of sodium and 121.4 g of chlorine

A 58.44 g sample of sodium chloride contains 22.99 g of sodium and 35.44 g of chlorine


Law of Multiple Proportions

• When two elements (call them A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small, whole numbers

63

John Dalton1766-1844


Oxides of Carbon

64

• Carbon combines with oxygen to form two different compounds, carbon monoxide and carbon dioxide

• Carbon monoxide contains 1.33 g of oxygen for every 1.00 g of carbon

• Carbon dioxide contains 2.67 g of oxygen for every 1.00 g of carbon

• Because there are twice as many oxygen atoms per carbon atom in carbon dioxide of in carbon monoxide, the oxygen mass ratio should be 2


Dalton’s Atomic Theory


• Dalton proposed a theory of matter based on it having ultimate, indivisible particles to explain these laws

1. Each element is composed of tiny, indestructible particles called atoms

2. All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements

3. Atoms combine in simple, whole-number ratios to form molecules of compounds

4. In a chemical reaction, atoms of one element cannot change into atoms of another element

they simply rearrange the way they are attached

Practice – Decide if each statement is correct according to Dalton’s model of the

atom• Copper atoms can combine with zinc

atoms to make gold atoms

• Water is composed of many identical molecules that have one oxygen atom and two hydrogen atoms


Practice – Decide if each statement is correct according to Dalton’s model of the atom

• Copper atoms can combine with zinc atoms to make gold atoms – incorrect; according to Dalton, atoms of one element cannot turn into atoms of another element by a chemical reaction. He knew this because if atoms could change it would change the total mass and violate the Law of Conservation of Mass.

• Water is composed of many identical molecules that have one oxygen atom and two hydrogen atoms – correct; according to Dalton, atoms combine together in compounds in small whole-number ratios, so that you could describe a compound by describing the number of atoms of each element in a molecule. He used this idea to explain why compounds obey the Law of Definite Proportions.


Practice – Decide if each statement is correct according to Dalton’s Model of the

Atom• Some carbon atoms weigh more than

other carbon atoms

• Because the mass ratio of Fe:O in wüsite is 1.5 times larger than the Fe:O ratio in hematite, there must be 1.5 Fe atoms in a unit of wüsite and 1 Fe atom in a unit of hematite


Practice – Decide if each statement is correct according to Dalton’s model of the

atom• Some carbon atoms weigh more than other carbon

atoms – incorrect; according to Dalton, all atoms of an element are identical.

• Because the mass ratio of Fe:O in wüsite is 1.5 times larger than the Fe:O ratio in hematite, there must be 1.5 Fe atoms in a unit of wüsite and 1 Fe atom in a unit of hematite – incorrect; according to Dalton, atoms must combine in small whole-number ratios. If you could combine fractions of atoms, that would mean the atom is breakable and Dalton’s first premise would be incorrect. You can get the Fe:Fe mass ratio to be 1.5 if the formula for wüsite is FeO and the formula for hematite is Fe2O3.


Cathode Ray Tube


• Glass tube containing metal electrodes from which almost all the air has been evacuated

• When connected to a high voltage power supply, a glowing area is seen emanating from the cathode

J.J. Thomson

• Believed that the cathode ray was composed of tiny particles with an electrical charge

• Designed an experiment to demonstrate that there were particles by measuring the amount of force it takes to deflect their path a given amount– like measuring the amount of force it takes to

make a car turn


72

Thomson’s Experiment

+++++++++++

-------------

Power Supply- +

Cathode Anode

Investigate the effect of placing an electric field around tube1. charged matter is attracted to an electric field2. light’s path is not deflected by an electric field

(-) (+)


Thomson’s Results• The cathode rays are made of tiny particles• These particles have a negative charge

– because the beam always deflected toward the + plate

• The amount of deflection was related to two factors, the charge and mass of the particles

• Every material tested contained these same particles • The charge:mass ratio of these particles was

−1.76 x 108 C/g– the charge/mass of the hydrogen ion is +9.58 x 104 C/g


Thomson’s Conclusions• If the particle has the same amount of charge

as a hydrogen ion, then it must have a mass almost 2000x smaller than hydrogen atoms!– later experiments by Millikan showed that the

particle did have the same amount of charge as the hydrogen ion

• The only way for this to be true is if these particles were pieces of atoms– apparently, the atom is not unbreakable


76

Millikan’s Oil Drop Experiment


Thomson’s Conclusions, cont’d

• Thomson believed that these particles were therefore the ultimate building blocks of matter– “We have in the cathode rays matter in a new state,

a state in which the subdivision of matter is carried very much further . . . a state in which all matter . . . is of one and the same kind; this matter being the substance from which all the chemical elements are built up.”

• These cathode ray particles became known as electrons


78

Electrons

• Electrons are tiny, negatively charged particles found in all atoms

• Cathode rays are made of streams of electrons

• The electron has a charge of −1.60 x 1019 C

• The electron has a mass of 9.1 x 10−28 g


A New Theory of the Atom

• Because the atom is no longer indivisible, Thomson must propose a new model of the atom to replace the first statement in Dalton’s Atomic Theory– rest of Dalton’s theory still valid at this point

• Thomson proposes that instead of being a hard, marble-like unbreakable sphere, the way Dalton described it, the atom actually had an inner structure


80

Thomson’s Plum Pudding Atom• The structure of the atom contains

many negatively charged electrons• These electrons are held in the

atom by their attraction for a positively charged electric field within the atom– there had to be a source of positive

charge because the atom is neutral– Thomson assumed there were no

positively charged pieces because none showed up in the cathode ray experiment


Predictions of the Plum Pudding Atom

• The mass of the atom is due to the mass of the electrons within it– electrons are the only particles in Plum

Pudding atoms, therefore the only source of mass

• The atom is mostly empty space– should not have a bunch of negatively

charged particles near each other as they would repel


82

Radioactivity• In the late 1800s, Henri Becquerel and

Marie Curie discovered that certain elements would constantly emit small, energetic particles and rays

• These energetic particles could penetrate matter

• Ernest Rutherford discovered that there were three different kinds of emissions– alpha, , rays made of particles with a mass

4x H atom and + charge– beta, , rays made of particles with a mass

~1/2000th H atom and – charge– gamma, , rays that are energy rays, not

particles

Marie Curie1867-1934


Rutherford’s Experiment• How can you prove something is empty space?

• Put something through it!– use large target atoms

• use very thin sheets of target so it will not absorb “bullet”

– use very small particle as bullet with very high energy • but not so small that electrons will affect it

• Bullet = alpha particles, target atoms = gold foil– particles have a mass of 4 amu & charge of +2 c.u.– gold has a mass of 197 amu & is very malleable


Rutherford’s Results

• Over 98% of the particles went straight through

• About 2% of the particles went through but were deflected by large angles

• About 0.005% of the particles bounced off the gold foil– “...as if you fired a 15” cannon shell at a piece

of tissue paper and it came back and hit you.”


Rutherford’s Conclusions

• Atom mostly empty space– because almost all the particles went straight

through

• Atom contains a dense particle that is small in volume compared to the atom but large in mass – because of the few particles that bounced back

• This dense particle is positively charged– because of the large deflections of some of the

particles


87

.

.

.

Nuclear Atom

•

••

•

• ••

• •••

•

•

•

• •

••

•

••

•

Plum PuddingAtom

If atom was likea plum pudding, all the particles

should go straight through

Almost all particles go straight through

Some particles go through, but are deflected due to

+:+ repulsion from the nucleus

A few of the particles

do not go through


88

Rutherford’s Interpretation – the Nuclear Model

1. The atom contains a tiny dense center called the nucleus

– the amount of space taken by the nucleus is only about 1/10 trillionth the volume of the atom

2. The nucleus has essentially the entire mass of the atom

– the electrons weigh so little they give practically no mass to the atom

3. The nucleus is positively charged – the amount of positive charge balances the negative

charge of the electrons

4. The electrons are dispersed in the empty space of the atom surrounding the nucleus


Structure of the Nucleus• Rutherford proposed that the nucleus

had a particle that had the same amount of charge as an electron but opposite sign – these particles are called protons– based on measurements of the nuclear

charge of the elements

• protons are subatomic particles found in the nucleus with a charge = +1.60 x 1019 C and a mass = 1.67262 x 10−24 g

• Because protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons


Relative Mass and Charge• It is sometimes easier to compare things to each other rather than

to an outside standard • When you do this, the scale of comparison is called a relative scale• We generally talk about the size of charge on atoms by comparing

it to the amount of charge on an electron, which we call −1 charge units– proton has a charge of +1 cu– protons and electrons have equal amounts of charge, but opposite signs

• We generally talk about the mass of atoms by comparing it to 1/12th the mass of a carbon atom with 6 protons and 6 neutrons, which we call 1 atomic mass unit– protons have a mass of 1 amu– electrons have a mass of 0.00055 amu, which is generally too small to be

relevant


Some Problems• How could beryllium have four protons stuck

together in the nucleus?– shouldn’t they repel each other?

• If a beryllium atom has four protons, then it should weigh 4 amu; but it actually weighs 9.01 amu! Where is the extra mass coming from?– each proton weighs 1 amu– remember, the electron’s mass is only about 0.00055

amu and Be has only four electrons – it can’t account for the extra 5 amu of mass


There Must Be Something Else!

• To answer these questions, Rutherford and Chadwick proposed that there was another particle in the nucleus – it is called a neutron

• Neutrons are subatomic particles with a mass = 1.67493 x 10−24 g and no charge, and are found in the nucleus1 amu

slightly heavier than a proton

no charge


Problems for Chapter 2

• 51-72, 75-78, 79-90

Ch 2. Atoms and Elements Atom Nucleus (Rutherford’s experiment) Electrons (Thomson’s experiment)

Documents

Transcript of Ch 2. Atoms and Elements Atom Nucleus (Rutherford’s experiment) Electrons (Thomson’s experiment)