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Atoms and moles

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Title: Atoms and moles


1
Atoms and moles
  • Chapter 3

2
Unit Essential Questions
  • 1) How have atoms been studied and understood
    throughout time?
  • 2) How do we handle the very small sizes of atoms
    in calculations?

3
Lesson Essential Question
  • 1) How do laws in chemistry support the existence
    of atoms?

4
Section 1 Substances Are Made of Atoms
  • First, a little history!
  • As early as 400 BC, a few people believed in
    atoms.
  • Democritus Greek philosopher, first to develop
    idea of the atom.
  • Had no evidence, only ideas.
  • Theory of the atom has changed over time to
    become what we have today.

5
Section 1 Cont.
  • Laws support atomic theory.
  • Developed from observations of compounds (how
    they are made up, how they react).
  • Recall what a compound is!
  • Atoms of two or more elements chemically
    combined.
  • Three laws were developed.

6
Introduction 1
  • It turns out that atoms and compounds are very
    similar to ingredients and food. How atoms come
    together to form compounds is similar to how
    ingredients come together to form food.
  • Think of a recipe for cooking or baking something
    and link this process to the law of conservation
    of mass.
  • The amount of ingredients you put into a recipe
    should equal the amount of food that comes out
    (but usually in a different form). No food is
    lost or gained (ideally)!

7
Law of Conservation of Mass
  • Antoine Lavoisier 1782
  • Mass cannot be created or destroyed in chemical
    or physical changes, only rearranged.
  • Example S O2 ? SO2
  • 32 g 32 g ? 64 g

8
Introduction 2
  • If you are baking a cake can you put together any
    amount/number of ingredients and always get out a
    cake?
  • No! So what can you say about amounts of
    ingredients needed?
  • The amounts are definite- only certain amounts of
    ingredients can make a cake.
  • Same is true for other recipes- given specific
    amounts of ingredients, only certain outcomes
    (food) are possible.
  • Atoms and compounds are the same! Atoms come
    together in definite amounts to form certain
    compounds.

9
Law of Definite Proportions
  • Atoms form compounds in specific, well defined
    proportions.
  • Proposed by Joseph Proust in 1797.
  • Example ethylene glycol (antifreeze)
  • Always 51 oxygen, 39 carbon and 10 hydrogen
  • C2H6O2
  • Table salt sodium chloride
  • 61 chlorine, 39 sodium
  • NaCl

10
Introduction 3
  • If you are making a hamburger at McDonalds, how
    many hamburgers and bun slices do you use?
  • One hamburger and two bun slices ratio is 12.
  • If you are making a Big Mac at McDonalds, how
    many hamburgers and bun slices do you use?
  • Two hamburgers and three bun slices ratio is
    23.
  • Based on this information, can hamburgers only be
    found as having one hamburger and two bun slices?
  • No! Other combinations are possible!
  • This is also true when atoms of certain elements
    combine
  • 2H2 O2 ? 2H2O H2 O2 ? H2O2

11
Law of Multiple Proportions
  • Two elements can combine to form two or more
    different compounds.
  • If the mass of the first element is held
    constant, the ratio of masses of the second
    element is always a small whole number ratio.
  • Example
  • What is the ratio by mass of O atoms?

Name of Compound Description Formula Mass O Mass N
nitrogen monoxide colorless gas NO 16.00 g 14.01 g
nitrogen dioxide poisonous brown gas NO2 32.00 g 14.01 g
12
12
Daltons Atomic Theory
  • 1808- John Dalton combined ideas of others to
    form an atomic theory.
  • Took Greek idea of atoms and turned it into a
    theory that could be tested.
  • Used the three laws previously discussed.
  • Was not 100 correct well look at new evidence
    that disproves some of his theory in Section 2.

13
Daltons 5 Principles of Atomic Theory
  1. All matter is composed of extremely small
    particles called atoms, which cannot be, created,
    destroyed, or subdivided.
  2. Atoms of a given element are identical in their
    physical and chemical properties.
  3. Atoms of different elements differ in their
    physical and chemical properties.
  4. Atoms of different elements combine in simple,
    whole number ratios to form compounds.
  5. In chemical reactions, atoms are combined,
    separated, or rearranged but never created or
    destroyed.

14
Daltons Atomic Theory- Which are Still True
Today?
  1. All matter is composed atoms, which cannot be
    subdivided, created, or destroyed.
  2. Atoms of a given element are identical in their
    physical and chemical properties.
  3. Atoms of different elements differ in their
    physical and chemical properties.
  4. Atoms of different elements combine in simple,
    whole number ratios to form compounds.
  5. In chemical reactions, atoms are combined,
    separated, or rearranged but never created or
    destroyed.

FALSE
FALSE
15
Lesson Essential Question
  • What are the components of the atom and why are
    they important?

16
Research Topics
  • Group 1 Thomsons discovery and how he
    discovered it
  • Group 2 Thomsons model of the atom- provide
    explanation picture
  • Group 3 Rutherfords discovery how he
    discovered it
  • Group 4 Rutherfords model of the atom- provide
    explanation picture
  • Group 5 Chadwicks discovery

17
Section 2 Structure of Atoms
  • Subatomic particles were discovered after
    Daltons theory.
  • The three we will discuss
  • Electron
  • Proton
  • Neutron
  • Others exist (quarks- make up neutrons and
    protons, leptons- make up electrons), but they
    are normally discussed in physics.

18
Electron
  • Discovered by JJ Thomson in the mid 1800s.
  • Studying electricity, not atoms, using a cathode
    ray tube.
  • Pumped all air out of glass tube.
  • Applied voltage to two metal plates, called
    electrodes.
  • Anode connected to positive terminal
  • Cathode connected to negative terminal
  • Glowing beam came out of cathode toward anode.
  • No atoms were inside the tube, so the beam must
    have come from the atoms in the cathode.
  • Beam was made of electrons.
  • Ray came from cathode- negatively charged.
  • Used in older TVs, computer monitors, and radar
    displays.

19
Further testing
  • Used an electric field in addition to the
    magnetic field to deflect the ray.
  • Further proof the beam was negatively charged.

20
What About Mass?
  • How could Thomson test the beam to see if it had
    mass? What can objects that have mass do that
    objects without mass cant do?
  • Move things!
  • So, Thomson placed a small paddle wheel in the
    beams path.
  • Wheel turned when hit by the beam. What does this
    tell you?
  • The beam consists of particles that have mass!

21
Further testing
  • Now that Thomson reached a conclusion, what
    should he do next?
  • Verify the results many times!
  • Retested using the same metal AND other metals.
    Both gave the same results.
  • Why was it important to test other metals and not
    just the same metal?
  • To show that electrons must exist in all atoms,
    not just the atoms in the first metal he tested.
  • Electrons- subatomic particles that have a
    negative charge.

22
Plum Pudding
  • How would you describe the locations of raisins
    in the plum pudding? Assume its the same inside.

23
Plum Pudding Model
  • Thomson proposed the plum-pudding model of the
    atom.
  • Electrons are embedded in a positive ball.
  • How did he know there should be a positive
    charge?

24
Further Searching
  • Imagine you were throwing a ball at a special
    wall, and 98 of the balls you threw at the wall
    went through it while the remaining 2 bounced
    off of it in various directions.
  • What can you conclude about the composition of
    this special wall?
  • Imagine the ball you were throwing had a positive
    charge on it. What would this tell you about any
    charges present in the wall?
  • Very similar to the gold foil experiment!

25
Searching for a Positive Subatomic Particle
  • Atoms are neutral, so positive particles must
    exist in atoms to balance out negative electrons.
  • Gold foil experiment conducted by Ernest
    Rutherford, student of Thomsons (1909).
  • Alpha particles (positive charge) directed at
    gold foil.
  • Most particles went through the foil.
  • Atoms must be mostly empty space.
  • Some particles were deflected.
  • There must be some concentrated positive area in
    atoms.

26
Nucleus Proton
  • Rutherford developed the idea of the nucleus.
  • Nucleus atoms positive central region,
    location of protons (and neutrons).
  • Protons- positive subatomic particle in the
    nucleus.
  • Mass is 2,000 times greater than an electron.
  • The nucleus is only 1/10,000 of the radius of the
    whole atom.
  • If the nucleus was the size of a marble, the
    entire atom would be the size of a football
    stadium!
  • Measure atoms radius in picometers (pm)
  • 10-12m.

27
Rutherfords Model of the Atom
  • Rutherfords experiments did not support
    Thomsons Plum Pudding model.
  • Developed the Planetary model- electrons look
    like planets orbiting the sun.
  • Lets visit an up-close picture of the gold foil.

28
Neutron
  • Discovered 30 years after the proton was found by
    James Chadwick in 1932.
  • Several people made observations about the
    neutron before Chadwick.
  • In studying a powerful beam, Chadwick was not
    able to deflect it with magnetic or electric
    fields.
  • Concluded the particles in the beam must be
    neutral in charge.
  • Neutrons subatomic particles found in the
    nucleus and have no electric charge.

29
Subatomic Particle Summary
Name Symbol Actual Charge (C) Common charge notation Mass (kg)
electron e- -1.602 x 10-19 -1 9.109 x 10-31
proton P 1.602 x 10-19 1 1.673 x 10-27
neutron n 0 0 1.675 x 10-27
30
Stability of Nuclei
  • How do protons stay together in the nucleus?
  • All protons are positively charged- why dont
    they push each other out of the nucleus?
  • Even though protons do repel one another in the
    nucleus, neutrons help hold them together.
  • Neutrons provide attractive forces without being
    subject to repulsive charge-based forces.

31
Atomic Number
  • Number of protons is unique to each element.
  • Can be used to identify elements.
  • Example atomic number 1 1 proton hydrogen
  • How many protons does carbon have?
  • What element has 80 protons?
  • The number of protons is the atomic number.
  • When atoms are neutral (no net charge), atomic
    number (number of protons) must equal number of
    electrons.
  • For a neutral atom p e- atomic

6
Hg (mercury)
32
Mass Number
  • Total number of subatomic particles in the
    nucleus.
  • Mass of n of p
  • Mass s are not unique (isotopes will explain
    later).
  • Can be used to find number of neutrons.
  • n mass - atomic
  • Examples
  • Hydrogen can have a mass 1, 2, or 3 but
    atomic number is always 1.
  • So, number of neutrons H atoms can have 0,1,2

Same as mass - p
33
Nuclear Symbols
  • Representation using symbols with numbers to
    identify the atomic number and/or mass number.
  • One method is name-mass .
  • Ex Carbon-12 or Carbon-14
  • Another method is called nuclear symbol.
  • Example ZX
  • A mass number Z atomic number X element
    symbol
  • 6C
  • Remember that the bottom number is always the
    same for any element, the top number can vary.
  • Because changing p changes the element!

A
12
34
Nuclear Symbols Ions
  • Not all atoms are neutral! Some have a charge.
  • Atoms with a charge are called an ion.
  • Ex O-2 and Na
  • The superscript gives the charge of the ion.
  • Only of e- are changed to produce ions. of p
    stays the same!
  • Positive charge e- lost
  • Subtract the charge number from the of e-
  • Negative charge e- gained
  • Add the charge number to the of e-

Na 11e- - 1e- 10e-
O-2 8e- 2e- 10e-
35
Isotopes
  • Isotope an atom with the same number of
    protons, but a different number of neutrons (and
    therefore a different mass ).
  • Identified using the two methods outlined on the
    previous slides.
  • Examples
  • 2He and 2He
  • copper-63 and copper-65

4
3
36
Lesson Essential Question
  • How can we describe the location of an electron
    in an atom?

37
Bohr Model
  • Developed after Rutherfords planetary model.
  • Recall that Rutherfords model showed that
    electrons orbit the positively charged nucleus
  • Problem why dont electrons crash into the
    nucleus?
  • Rutherfords model could not answer this question.

38
Bohr Model
  • Bohr proposed that electrons can only orbit the
    nucleus at certain energy levels- they cannot
    exist anywhere in between.
  • Similar to the rungs on a ladder- you cant step
    in between them, theres nowhere to put your foot
    down!
  • Thus, electrons dont crash into the positive
    nucleus because there is no energy level (orbit)
    for them to exist in.

39
Electrons and Light
  • Einstein (1905) proposed light had properties of
    particles in addition to wave properties.
  • Photoelectric effect a certain amount of energy
    is needed to remove an electron from a piece of
    metal when struck with light.
  • If light only acted as waves, any frequency would
    eventually have enough energy to remove an
    electron. But this was not seen!
  • Only certain frequencies with certain energies
    could remove an electron.

40
Electrons as Particles and Waves
  • DeBroglie (1924) pointed out that electrons act
    as waves as well as particles.
  • Quantum model of the atom used orbitals regions
    where electrons are likely to be found.
  • Also called
  • electron clouds.
  • No sharp boundaries.
  • Uses probability.

41
Match the model to the scientist/theory
  • Thomson (Plum Pudding)
  • Bohr
  • Rutherford (Planetary)
  • Quantum

42
Light Emission
  • Electrons have a certain energy level where they
    are located ground state low energy.
  • Higher state excited state higher energy.
  • When they are removed or moved their energy
    changes.
  • The difference in energy is usually released as
    light.
  • Each element can give a unique line-emission
    spectrum.
  • Bohr developed an equation to calculate the
    energy of each electron.
  • This led to many people accepting his model of
    the atom.

43
Electron Excitation
44
Electromagnetic Spectrum
45
Warm-Up Question
  • What is the purpose of an address? Why is each
    component necessary?
  • Examine the following address
  • 430 New Schaefferstown Road
  • Bernville, PA
  • Order the components of the address from most
    general to most specific.
  • PA, Bernville, New Schaefferstown Rd., 430

46
Quantum Numbers
  • Like an address, electrons can be identified by
    where they reside in an atom.
  • The 4 parts to this address are called quantum
    numbers.
  • Each quantum number further pinpoints an
    electron.
  • In other words, quantum numbers separate
    electrons from one another- they let you tell
    them apart.

47
Principal Quantum Number
  • Principal quantum , n, tells the energy level.
  • n can only be positive integers (1, 2, 3, )
  • The larger the n value, the farther the e- is
    from the nucleus, and the greater its energy.
  • In terms of an address, this would be like the
    state- it gives you the most general idea of
    where an electron is.

n 1 (first energy level)
n 2 (second energy level)
48
Angular Momentum Quantum Number
  • Angular momentum quantum , l, tells the
    sublevel/orbital.
  • Can be zero or any positive whole number.
  • Each sublevel has a different shape.
  • Letters designate shapes of different l values
  • l 0 s
  • l 1 p
  • l 2 d
  • l 3 f

p
d
s
49
Angular Momentum Quantum Number Continued
  • In terms of an address, this would be like the
    city.
  • Cities give us a more specific area to look for
    someone within a state.
  • Sublevels further specify where an electron is
    within an energy level.

50
Magnetic Quantum Number
  • Magnetic quantum , m, tells the orbitals
    orientation in space.
  • Each orbital shape can have different numbers of
    orientations.
  • More orientations more orbitals present in a
    sublevel.
  • s has 1 orbital because it only has 1 possible
    orientation.
  • p has 3 orbitals because it has 3 possible
    orientations.
  • d has 5 orbitals, and f has 7.
  • Notice the pattern- the number of orbitals (or
    orientations) increases by 2 for each sublevel.

51
A Further Look at Orientations
52
Magnetic Quantum Number
  • In terms of an address, this would be like the
    street.
  • The street that a person lives on allows us to
    further isolate where that person is located
    within a city.
  • Orientations of orbitals further specify where an
    electron is within a sublevel.

53
Spin Quantum Number
  • Spin quantum number, s, tells the electrons
    orientation within an orbital.
  • Can only have 2 values, which are symbolized in 3
    ways
  • 1/2 and -1/2
  • and
  • clockwise and counterclockwise
  • Note any orbital can only hold up to TWO
    electrons!
  • In an address, this would be like the street
    number.
  • Once we have the street number of a person, we
    know exactly where to find them.
  • The same goes for an electron when we know the
    spin.

54
Another Look at Spin
  • The up and down arrows will come into play when
    we learn to write electron configurations.

55
Thinking Ahead- Electron Configurations
  • Can 2 e- in the same orbital have the same spin
  • ( or )? (This would be like two homes
    on the same street having the same house number.)
    Why or why not?
  • Do you think e- would prefer to be closer to the
    nucleus at a lower energy level or farther away
    from the nucleus at a higher energy level?
  • Hint think about charges and energy involved.

56
Thinking Ahead- Electron Configurations
  • If given the choice, do you think 2 e- would
    rather be paired together in the same orbital or
    be alone in different orbitals?
  • Hint think about charges.

57
Electron Configuration Rules
  • When determining the placement of electrons in an
    atom, three rules must be followed.
  • Pauli Exclusion Principle no two electrons in
    the same atom can have the same four quantum
    numbers.
  • Just like no two places can have the same
    address!
  • aufbau principle electrons fill up the lowest
    energy level first.

58
Electron Configuration Rules Cont.
  • Hunds rule One electron must occupy each
    orbital before pairing.
  • Electrons are negatively charged and repel each
    other, so they spread out as much as possible.
  • Think of this as the movie theater rule.

59
Electron Configurations
Row numbers energy levels (n) Blocks shapes
1
2
3
s
p
4
d
5
(n)
(n)
6
(n-1)
7
f
(n-2)
Remember atomic e- if theres no charge!
60
Types of Electron Configurations
  • Full
  • All electrons are written out.
  • Example Write the electron configuration for an
    atom of N.
  • First, determine the atomic number.
  • This tells you how many electrons N has.
  • Then write the electron configuration using the
    periodic table.
  • 1s22s22p3
  • Check yourself! The superscripts should add up to
    the e- (atomic ) 223 7

7
61
Types of Electron Configurations
  • Abbreviated Cont.
  • Find the noble gas youll need to use
  • (1) Go up one row from the element youre
    writing the abbreviated configuration for.
  • (2) Go all the way to the right-most column on
    the periodic table. This element is the noble gas
    youll use. Write the symbol for this element in
    brackets .
  • Determine how many electrons the noble gas has,
    and count up to this many using the orbital
    filling diagram. Where you end is implied by the
    noble gas in brackets.
  • Pick up with your configuration as you normally
    would.

62
Types of Electron Configurations
  • Orbital Diagram
  • May use full or abbreviated AND involves the use
    of horizontal lines (___) for each orbital.
  • Arrows are used with each line to show electrons.
  • Recall that up ( ) and down ( ) arrows are used
    to show different spins! (spin quantum number
    s).

63
Lesson Essential Question
  • How do we count and work with large quantities of
    atoms?

64
Section 4 Counting Atoms
  • Atomic mass units created just to measure masses
    of atoms because theyre so small.
  • Abbreviated amu.
  • Daltons (Da) can also be used.
  • Use the values on the periodic table.

65
The Mole Avogadros Number
  • Mole SI base unit amount of substance.
  • Like a counting unit.
  • 1 dozen 12 eggs
  • Number of particles in one mole is called
    Avogadros number.
  • 1mole 6.022 x 1023 particles
  • Extremely large 602,200,000,000,000,000,000,000!
  • Atoms, molecules, etc. can be used for labels
    instead of particles.
  • Named after work that Amadeo Avogadro completed.

66
Example
  • If you have a dozen people and a dozen cars, what
    do they have in common?
  • They both have 12!
  • Since the numbers are the same does that mean
    that their masses are the same?
  • No! The heavier items will have a greater mass!
  • This is the same for atoms! If you have a mole of
    carbon and a mole of gold they both have
    6.022 x 1023 atoms. But a gold atom is heavier
    than a carbon atom, so the mole of gold has a
    larger mass!

67
The Mole Molar Mass
  • To convert between moles and grams, we use molar
    mass (conversion factor).
  • Molar mass mass in grams of one mole of an
    element.
  • Units grams/mole or g/mol
  • Example Cu 63.55 amu 63.55 g/mol
  • We will round all atomic masses or molar masses
    to 2 places after the decimal point.
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