topic 2.1: atomic structure honors chemistry 2014-15 mrs. peters 1

Topic 2.1: Atomic Topic 2.1: Atomic StructureStructure

Honors Chemistry 2014-15Mrs. Peters

1

Atomic Structure Atomic Structure

2.1: The nuclear atom2.1: The nuclear atom

EI: The mass of the atoms is concentrated in its minute, positively charged nucleus.

NOS: 1.Evidence and improvements in instrumentation – alpha particles were used in the development of the nuclear model of the atom that was first proposed by Rutherford. (1.8)2.Paradigm shifts- the subatomic particle theory of matter represents a paradigm shift in science that occurred in the late 1800s (2.3)

2



Understandings:1.Atoms contain a positively charged dense nucleus composed of protons and neutrons (nucleons)2.Negatively charged electrons occupy the space outside the nucleus3.The mass spectrometer is used to determine the relative atomic mass of an element from its isotopic composition.

3



Applications and Skills:1.Use of the nuclear symbol notation A

ZX to deduce the number of protons, neutrons, and electrons in atoms and ions.2.Calculations involving non-integer relative atomic masses and abundance of isotopes from given data, including mass spectra.

4

NOS: Paradigm ShiftNOS: Paradigm Shift

History behind Atomic Theory• Democritus (420 BCE) first proposed the idea

that matter may be made up of small, indivisible particles called atoms.

• Aristotle (384-322 BCE) Greek philosopher; matter composed of earth, air, fire, water. This view dominated thought until 17th century

5


History behind Atomic Theory• Atomism developed in Chinese & Arabic cultures

during the Dark Ages in Europe.

• John Dalton (1766-1844) was the first to base atomic theory on scientific evidence.

6


Dalton’s Atomic Theory• Elements are made of tiny particles called atoms.

• All atoms of a given element are identical. The atoms of a given element are different from those of any other element.

7


Dalton’s Atomic Theory• Atoms of one element can combine with atoms of

other elements to form compounds. A given compound always has the same relative number of types of atoms.

• Atoms cannot be created, nor divided into smaller particles, nor destroyed in the chemical process. A chemical reaction simply changes the way atoms are grouped together.

8


Evidence for sub-atomic particles

1897: J.J. Thomsen: Cathode Ray TubeEvidence for electrons: Bent a stream of rays originating from the negative electrode (cathode). Stream of particles with mass & negative charge.

9



1909: Ernest Rutherford: Gold FoilEvidence for protons & nucleus: Alpha particles deflected passing through gold foil

10



1932: James Chadwick: BerylliumEvidence for neutrons: Alpha particles caused beryllium to emit rays that could pass through lead but not be deflected,

11

U1. and U2. Atomic U1. and U2. Atomic StructureStructure

Sub-Atomic Particles: Proton: Located in the nucleus

Relative charge of +1

Relative mass of 1 amu

Neutron: Located in the nucleusRelative charge of 0 Relative mass of 1

amu

12www.green-planet-solar-energy.com, 3.bp.blogspot.com


Sub-Atomic ParticlesElectron: Located in cloud

surrounding the nucleus Relative charge of –1Relative mass of

0.0005 amu

13

www.green-planet-solar-energy.com, 3.bp.blogspot.com


Nucleus consists of protons and neutrons with theelectrons surrounding the nucleus.

In a neutral atom, the #protons = # electrons.

14

A1. Nuclear Symbol NotationA1. Nuclear Symbol Notation

Atomic Number (Z)

The atomic number is the number of protons in the nucleus. It determines the identity of an atom.

• All oxygen atoms have 8 protons in the nucleus• All lead atoms have 82 protons in the nucleus

15


Atomic Number (Z)

It also tells us the number of electrons in a neutral atom

• A neutral sodium atom contains 11 protons and 11 electrons

• A neutral bromine atom contains 35 protons and 35 electrons

16


Mass Number (A)

It is not practical to measure the masses of atoms in grams

due to their small size. Scientists devised a measurement

called atomic mass units (amu).

17


Mass Number (A)

• Protons have a mass of 1 amu• Neutrons have mass of 1 amu• Electrons have mass of 0 amu.

Mass Number of atoms = # protons + # neutrons

18


Mass Number (A)Mass Number of atoms = # protons + #

neutrons

**Round the Relative Atomic Mass to a whole number to find the Mass Number

Ex:• Lithium = 6.94 Mass Number is 7• Magnesium = 24.31 Mass Number is 24

19

How to read the How to read the Periodic TablePeriodic Table

20

Lithium3

Li6.94

Element Name

Atomic Number

Element Symbol

Relative Atomic Mass


Atomic Name: Element Name - A (mass number)

Ex: Carbon-12

Nuclear Symbol:

21

XA

ZElement Symbol

Mass Number (Protons + Neutrons)

Atomic Number

Important TermsImportant Terms

IsotopeAtoms of the same element can have different

numbers of neutrons, thus they will have different atomic

masses. These are called isotopes of the element.

These are the same element, just different numbers of neutrons and mass.

22


Isotope Example

There are three isotopes of hydrogen:• Hydrogen-1 has 1 proton, 1 electron, 0

neutrons• Hydrogen-2 has 1 proton, 1 electron, 1 neutron• Hydrogen-3 has 1 proton, 1 electron, 2

neutrons

23

A1. Deduce the symbol given its mass A1. Deduce the symbol given its mass number and atomic numbernumber and atomic number

• Consider an atom that has an atomic number of 29 and a mass number of 63. What is its name and symbol?

Name:Symbol:

24

A. Deduce the symbol given its mass A. Deduce the symbol given its mass number and atomic numbernumber and atomic number

• Consider an atom that has an atomic number of 29 and a mass number of 63. What is its name and symbol? atomic number of 29 identifies it as copper

Name: Copper-63 Symbol: 63 Cu

29

25


•Consider an atom that has A=32 and Z=16. What is its name and symbol?

Name: Symbol:

26


• Consider an atom that has A=32 and Z=16. What is its name and symbol?Z=16 identifies it as sulfur

Name: Sulfur-32Symbol: 32 S

16

27


• Consider an atom that has an atomic number of 74 and a mass number of 185. What is its name and symbol?

• Consider an atom that has A=127 and Z=53. What is its name and symbol?

28


• Consider an atom that has an atomic number of 74 and a mass number of 185. What is its name and symbol? atomic number of 74 identifies it as tungsten

Name: Tungsten-185 Symbol: 185 W

74

• Consider an atom that has A=127 and Z=53. What is its name and symbol?Z=53 identifies it as iodine

Name: Iodine-127Symbol: 127 I

53

29

A1. Deduce protons, neutrons, and A1. Deduce protons, neutrons, and electrons in atoms and ions from the A, Z, electrons in atoms and ions from the A, Z,

and chargeand charge

• Consider the neutral carbon-12 atom. Find the A, Z, protons, neutrons, electrons, and symbolName is Carbon-12 Atomic mass (A) = 12Atomic number (Z) = 6Protons = 6 (atomic number)Neutrons = 6 (mass – protons)Electrons = 6 (neutral atom so same as protons)

Symbol is 12C6

30



• Consider an atom that has 9 protons, 9 electrons, and 10 neutrons. What is its atomic number, atomic mass, name, and symbol?

Z=9 (atomic number = # protons) A=19 (atomic mass = protons + neutrons)Fluorine-19 (name and mass) 19F (neutral because protons = electrons)

9

31



• Consider a neutral atom with A=75 and Z=33. How many protons, neutrons, and electrons are in the atom. What is the name and symbol?


32

A1. Deduce protons, neutrons, and A1. Deduce protons, neutrons, and electrons in atoms and ions from A, Z, and electrons in atoms and ions from A, Z, and

chargecharge


Protons = 33 Neutrons = 42 Electrons = 33

Name: Arsenic-75

Symbol: 75As

33


Protons = 33 Neutrons = 44 Electrons = 33

Name: Arsenic-77

Symbol: 77As

33

33


Ions are charged particles formed when atoms gain or loseelectrons resulting in unequal numbers of protons and

electrons

Cations: Atoms that lose electrons become positively charged

Anions: Atoms that gain electrons become negatively charged

34

A1. Use of Nuclear Symbol A1. Use of Nuclear Symbol NotationNotation

Nuclear Symbol:

35

XA

ZElement Symbol

Mass Number (Protons + Neutrons)

Atomic Number

+Charge (+, - or nothing)Determined by electrons


• How many protons, neutrons, and electrons are in an ion of K-39 that has lost one electron? What is the charge of the ion? What is its symbol?Protons = 19 Neutrons = 20 Electrons = 18Charge = 1+ or +1 Symbol is 39K1+

19

36

A1. Deduce from nuclear symbol A1. Deduce from nuclear symbol

notationnotation

• The symbol of an anion is 31P 3- . Calculate the number of

15

protons, neutrons, and electrons. What is Z and what is A?

• What is the symbol of a species containing 26 protons, 30 neutrons, and 23 electrons?

• What is the symbol of a species with A=56, Z=26, and 24 electrons?

37

A1. Deduce the number of protons, A1. Deduce the number of protons, neutrons, and electrons in atoms and ions neutrons, and electrons in atoms and ions

from A, Z, and chargefrom A, Z, and charge

• The symbol of an anion is 31P 3- . Calculate the number

15

protons, neutrons, and electrons. What is Z and what is A? #P = 15; #N = 16; #E = 18; Z= 15; A = 31

• What is the symbol of a species containing 26 protons, 30 neutrons, and 23 electrons?

56Fe 3+

26

• What is the symbol of a species with A=56, Z=26, and 24 electrons?

56Fe 2+

26

38

A2. Discuss the use of radioisotopes.A2. Discuss the use of radioisotopes.

Radioisotopes: isotopes of elements that have become radioactive because the nucleus is unstable and breaks down spontaneously emitting radiation.

Radioisotopes can occur naturally or be created artificially

Examples: Carbon-14; Iodine-125; Strontium-90; Cobalt-60, Iodine-131

39

A2. Discuss the use of radioisotopes.A2. Discuss the use of radioisotopes.

Uses of Radioisotopes

• Nuclear power generation• Sterilization of surgical instruments• Crime detection• Food preservation• Dating artifacts• Treating and diagnosing disease

40

U3. The Mass SpectrometerU3. The Mass Spectrometer

Mass Spectrometers:Instruments that measure charge-to-mass ratio of charged particles.Used to measure masses of isotopes as well as isotopic abundance

41


How a Mass Spec Works:1. Vaporization: sample is heated to

gas state2. Ionization: sample gas is turned

into ions by blasting free electrons to knock electrons off from the gas atoms, creating positive ions

3. Acceleration: increases the speed of particles, using an electric field

42


How a Mass Spec Works:4. Deflection: using an

electromagnet to create a magnetic field, amount of deflection depends on mass and charge of the ion (think of cars going around a corner)

5. Detection: measures both mass and relative amounts (abundance) of all the ions present

43


• Mass Spectrometer Video

• http://www.youtube.com/watch?feature=fvwp&v=lxAfw1rftIA&NR=1

44

U3. Relative Atomic MassU3. Relative Atomic Mass


Mass numbers (atomic mass) on the periodic table are weighted averages of the isotopes.

Based on 12C. Has 6 protons, 6 neutrons, and 6 electronsHas a relative atomic mass of exactly 12.000

45



One amu is exactly 1/12 of the mass of a carbon-12 atom.

All other isotopes are measured compared to this value.

46


Average relative atomic mass: the weighted average for all of

the isotopes of a given element, based on the percent abundance of each

47


To determine Average Relative Atomic Mass:

• Need masses of each isotopes

• Need abundance (percentage) of each isotopeo The mass spec is used to

determine these values

• This is the value shown on the periodic table

48


• A sample of neon is placed in the mass spectrometer

49


• A sample of neon is placed in the mass spectrometer

• The results show the abundance for each isotope of an elemento 90.92% is neon-20o 0.26% is neon-21o 8.82% is neon-22

50

A2. Calculate non-integer relative atomic A2. Calculate non-integer relative atomic masses and abundance of isotopes from masses and abundance of isotopes from

given data.given data.

Example1.95.5% = .955 .5%=0.005

2.24.5 x .955= 23.4 and 23.7 x .005 = .119

3.23.4 + .119 = 23.5194.23.5 amu

How to Determine Relative Atomic Mass

1. Convert the percent abundance for each isotope into decimal2. Multiply the mass for each isotope by the abundance 3. Add all product values from step 2. 4. Include amu for the units of the value.

51



Three isotopes of magnesium occur in nature.

Their abundances and masses, determined by

mass spectrometry, are listed in the table on

the right. Use this information to calculate the

atomic weight of magnesium.

• Three isotopes: 24, 25, 26• Percentage of each isotope: Given• Multiply the percent of each isotope by

its mass23.98504 x .7899 = 18.95 amu 24.98584 x .1000 = 2.499 amu25.98259 x .1101 = 2.861 amu

• Add these values = 24.31 amu

Isotope % Abundance

Mass (amu)

24Mg 78.99 23.98504

25Mg 10.00 24.98584

26Mg 11.01 25.98259

52



Calculate the atomic weight of chromium using the following data for the

percent natural abundance and mass of each isotope:4.35% 50Cr (49.9461 amu); 83.79% 52Cr (51.9405 amu);

9.50% 53Cr (52.9406 amu); 2.36% 54Cr (53.9389 amu)

53

A2. Calculate non-integer relative A2. Calculate non-integer relative atomic masses and abundance of atomic masses and abundance of

isotopes from given data.isotopes from given data.

Calculate the atomic weight of chromium using the following data for the

percent natural abundance and mass of each isotope:4.35% 50Cr (49.9461 amu); 83.79% 52Cr (51.9405 amu);

9.50% 53Cr (52.9406 amu); 2.36% 54Cr (53.9389 amu)

49.9461 x .0435 = 2.17 amu51.9405 x .8379 = 43.52 amu52.9406 x .0950 = 5.03 amu53.9389 x .0236 =+ 1.27 amu

51.99 amu

54



Determine the atomic weight of lead using

the data from the mass spectrum of lead

• Four isotopes: 204, 206, 207, 208• Percentage of each isotope:

Total # isotopes is 10 (1+2+2+5)204: 1/10 = 10% 206: 2/10 = 20%207: 2/10 = 20% 208: 5/10 = 50 %

• Multiply the percent of each isotope by its mass204 x .1 = 20.4 206 x .2 = 41.2207 x .2 = 41.4 208 x .5 = 104

• Add these values20.4 + 41.2 + 41.4 + 104 = 207

55

Mass Spectrum of Lead

A2. Calculate non-integer relative A2. Calculate non-integer relative atomic masses and abundance of atomic masses and abundance of

isotopes from given data.isotopes from given data.

The atomic weight of gallium is 69.72 amu. The masses of the naturally occurring isotopes are 68.9257 amu for 69Ga and

70.9249 amu for 71Ga. Calculate the percent abundance of each isotope.

• Let x = % abundance of 69Ga. Then 1-x = % abundance of 71Ga.

• 68.9257x + 70.9249(1-x) = 69.72 amu68.9257x + 70.9249 – 70.9249x = 69.72-1.9992x = -1.20x = 0.600 = decimal value of 69Ga so 60.0% 69Ga1-x = 0.400 = decimal value 71Ga so 40.0 % 71Ga

56



The atomic weight of copper is 63.546 amu. The masses of the two naturally occurring isotopes are 62.9298 amu for 63Cu and 64.9278 amu for 65Cu. Calculate the percent of 63Cu in naturally occurring

copper.

57



The atomic weight of copper is 63.546 amu. The masses of the two naturally occurring isotopes are 62.9298 amu for 63Cu and 64.9278 amu for 65Cu. Calculate the percent of 63Cu in naturally occurring

copper.

• Let x = % abundance of 63Cu. Then 1-x = % abundance of 65Cu.

• 62.9298x + 64.9278(1-x) = 63.546 amu62.9298x + 64.9278 – 64.9278x = 63.546-1.998x = -1.382x = 0.6917 = decimal of 63Cu so 69.17% 63Cu

58

topic 2.1: atomic structure honors chemistry 2014-15 mrs. peters 1

Documents

paradigm shift evidence

atomic theory atomism

atomic theory democritus

relative mass

stream of particles

way atoms

paradigm shift history

integer relative atomic