ATOMS AND MOLES

Chapter 3

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?

Lesson Essential Question:

1) How do laws in chemistry support the existence of atoms?

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.

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.

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)!

Law of Conservation of Mass

 Antoine Lavoisier 1782  Mass cannot be created or destroyed in chemical or physical changes, only rearranged.

 Example: S + O 2  32 g + 32 g   SO 64 g 2

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.

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  C 2 H 6 O 2  Table salt = sodium chloride  61% chlorine, 39% sodium  NaCl

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 1:2.

 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 2:3.

 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: 2H 2 + O 2  2H 2 O H 2 + O 2  H 2 O 2

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:

Name of Compound Description Formula Mass O Mass N nitrogen monoxide nitrogen dioxide colorless gas poisonous brown gas NO NO 2 16.00 g 32.00 g 14.01 g 14.01 g

 What is the ratio by mass of O atoms?

1:2

Dalton

’

s 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 – we ’ ll look at new evidence that disproves some of his theory in Section 2.

Dalton

’

s 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.

Dalton

’

s Atomic Theory- Which are Still True Today?

1.

2.

3.

4.

5.

All matter is composed atoms, which cannot be subdivided, created, or destroyed.

FALSE

Atoms of a given element are identical in their physical and chemical properties.

FALSE

Atoms of different elements differ in their physical and chemical properties.

Atoms of different elements combine in simple, whole number ratios to form compounds.

In chemical reactions, atoms are combined, separated, or rearranged but never created or destroyed.

Lesson Essential Question:

What are the components of the atom and why are they important?

Research Topics

 Group 1: Thomson’s discovery and how he discovered it  Group 2: Thomson’s model of the atom- provide explanation & picture  Group 3: Rutherford’s discovery & how he discovered it  Group 4: Rutherford’s model of the atom- provide explanation & picture  Group 5: Chadwick’s discovery

Section 2: Structure of Atoms

 Subatomic particles were discovered after Dalton ’ s 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.

Electron

 Discovered by JJ Thomson in the mid 1800 ’ s.

 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.

Further testing…

 Used an electric field in addition to the magnetic field to deflect the ray.

 Further proof the beam was negatively charged.

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 can’t do?

 Move things!

 So, Thomson placed a small paddle wheel in the beam ’ s path.

 Wheel turned when hit by the beam. What does this tell you?

 The beam consists of particles that have mass!

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.

Plum Pudding

 How would you describe the locations of raisins in the plum pudding? Assume it’s the same inside.

Plum Pudding Model

  Thomson proposed the “ plum-pudding ” of the atom.

model  Electrons are embedded in a positive ball.

How did he know there should be a positive charge?

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!

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 Thomson ’ s (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.

Nucleus & Proton

 Rutherford developed the idea of the nucleus.

 Nucleus – atom ’ s 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 atom ’ s radius in picometers (pm) = 10 -12 m.

Rutherford

’

s Model of the Atom

 Rutherford ’ s experiments did not support Thomson ’ s Plum Pudding model.

 Developed the Planetary model- electrons look like planets orbiting the sun.

 Let’s visit an up-close picture of the gold foil.

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.

Subatomic Particle Summary

Name

electron

Symbol

e -

Actual Charge (C) Common charge notation

-1.602 x 10 -19 -1

Mass (kg)

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

Stability of Nuclei

 How do protons stay together in the nucleus?  All protons are positively charged- why don ’ t 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.

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?

6

 What element has 80 protons?

Hg (mercury)

 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 #

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 # Same as: mass # - #p +  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

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:  A Z X A = mass number ; Z = atomic number; X = element symbol  12 6 C  Remember that the bottom number is always the same for any element, the top number can vary.

 Because changing #p + changes the element!

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 p + are changed to produce ions. # of stays the same!

 Positive charge = e

-

lost Na + : 11e

-

- 1e

-

= 10e

-

 Subtract the charge number from the # of e

-

 Negative charge = e

-

gained O -2 : 8e

-

 Add the charge number to the # of e

-

+ 2e

-

= 10e

-

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:  4 2 He and 3 2 He  copper-63 and copper-65

Lesson Essential Question:

How can we describe the location of an electron in an atom?

Bohr Model

 Developed after Rutherford ’ s planetary model.

 Recall that Rutherford ’ s model showed that electrons orbit the positively charged nucleus:  Problem: why don ’ t electrons crash into the nucleus?

 Rutherford ’ s model could not answer this question.

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 can ’ t step in between them, there ’ s nowhere to put your foot down!

 Thus, electrons don ’ t crash into the positive nucleus because there is no energy level (orbit) for them to exist in.

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.

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.

Match the model to the scientist/theory

 Thomson (Plum Pudding)  Bohr  Rutherford (Planetary)  Quantum

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.

Electron Excitation

Electromagnetic Spectrum

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

Quantum Numbers

 Like an address, electrons can be identified by where they “ reside ” in an atom.

 The 4 parts to this ‘ address ’ numbers.

are called quantum  Each quantum number further pinpoints an electron.

 In other words, quantum numbers separate electrons from one another- they let you tell them apart.

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

-

nucleus, and the greater its energy.

is from the  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)

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 s p d

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.

Magnetic Quantum Number

 Magnetic quantum #, m, tells the orbital ’

s

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.

A Further Look at Orientations

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.

Spin Quantum Number

 Spin quantum number, s, tells the electron ’

s

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.

Another Look at Spin

 The up and down arrows will come into play when we learn to write electron configurations.

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.

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.

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.

Electron Configuration Rules Cont.

 Hund ’ s 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.

Electron Configurations

1 2 7 3 4 5 6 s

(n)

Row numbers = energy levels (n) Blocks = shapes d

(n-1)

p

(n)

f

(n-2)

Remember: atomic # = #e if there ’s no charge!

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.

 1s 2 2s 2 2p 3  Check yourself! The superscripts should add up to the # e (atomic #): 2+2+3 = 7

Types of Electron Configurations



Abbreviated Cont.

 Find the noble gas you’ll need to use: (1) Go up one row from the element you’re 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 you’ll 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.

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 ’ ).

Lesson Essential Question:

How do we count and work with large quantities of atoms?

Section 4: Counting Atoms

 Atomic mass units: created just to measure masses of atoms because they ’ re so small.

 Abbreviated amu.

 Daltons (Da) can also be used.

 Use the values on the periodic table.

The Mole & Avogadro

’

s Number

  Mole – SI base unit – amount of substance.

 Like a counting unit.

 1 dozen = 12 eggs Number of particles in one mole is called Avogadro ’

s

number.

 1mole = 6.022 x 10 23 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.

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 10 larger mass!

23 atoms. But a gold atom is heavier than a carbon atom, so the mole of gold has a

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.