CHEMISTRY NOTES: Unit 2: Structure of Matter

NAME: ______PERIOD: ______

CHEMISTRY NOTES: Unit 2: Structure of Matter

Content Outline: History of the Atomic Model (2.1)

1. Democritus (400 B. C.)

1. He was a Greek philosopher of science.
2. First to use the term “______” to describe the basic particle of nature.

1. “atom” means “indivisible”
2. Atom –

1. John Dalton (1808)

1. He was an English schoolteacher.
2. He was the first to propose an “______” that contains the 5 following statements:

1. All matter is composed of extremely small particles called “______”.
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-number ratios to form chemical

compounds.

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

1. Joseph John Thomson (1897)

1. He was an English Physicist.
2. He worked with glass gas-filled tubes referred to as ______.

1. The glass tubes were filled with a gaseous element under low pressure.
2. He then passed an electrical current through the gas using a battery and wires.

1. The electrical current caused the gas within the tube to intensely glow with a beam (“ray”).

1. Magnets could make the “ray” move/deflect in various directions.
2. The ray is being deflected by the ______ of the magnet.
3. Negative charge repels/deflects like negative charges.
4. The ray is made of a negative charge that Thompson called ______ (since they were associated with the electrical current.)

1. The electrical current came into the chamber (by a wire) at the ______end. (The end where electricity enters the tube.)
2. Hence the term Cathode Ray tubes.

1. Further investigations using different elements in Cathode-Ray tubes confirmed that every element’s atoms possess ______.
2. He proposed the “______” model of atoms.

1. It stated that ______ charged electrons are evenly placed inside a ______- charged mass.

1. Robert A. Milliken (1909)

1. He was an American Physicist.
2. He was the first to measure ______-.

The symbol for an electron is: ______

1. Electron charge = 1.602 x 10-19 Coulombs.

1. This is an extremely small quantity of energy.

1. Electron mass = 9.11 x 10-31 kg

1. Electrons are 1/1837th the mass of a single proton or neutron.

1. Milliken’s experiments allow for 2 inferences (conclusions based upon evidence and reasoning) to be made:

1. Because atoms, in the natural state, are ______, they must also contain an equal amount of ______.
2. Because electrons have so little mass, atoms must contain other particles with much greater mass (protons & neutrons).

1. Ernest Rutherford, Hans Geiger, and Ernest Marsden (1911)

1. Geiger and Marsden were students of Rutherford a New Zealand Physicist.
2. They performed the ______-.
3. They used high-energy alpha particle radiation (2 protons & 2 neutrons ejected from a decomposing, radioactive element) to bombard a piece of gold foil that was surrounded by a fluorescent screen.

1. As alpha particles struck the fluorescent screen, they would produce a small detectable burst of light.
2. As the experiment was running, they detected light burst mainly behind the gold foil, but also occasionally all around the ring.

1. These bursts of light around the ring were because of the positively charged alpha particles been deflected by positively charged particles in the atoms of the foil.
2. The particles became known as ______.

1. Just as with the electrons, positive charges repel/deflect like positive charges.

1. As most of the bursts of light occurred behind the gold foil, they concluded that the majority of space in an atom is “empty space” that the alpha particles travelled through and never hit anything.

1. ______proposes the idea of the neutrally charged ______ particle in 1920.

1. Niels Bohr (1913)

1. He was also a student of Rutherford’s.
2. He proposed the ______ of an atom.

1. The electrons move in a circular pattern around the positively charged center. (Much like the planets revolve around the sun.)

1. Dmitri Ivanenko & Victor Ambartsumian (1930)

1. These gentlemen were Russian Physicists.
2. They proposed a model of the nucleus of an atom that is composed of positively charged ______ and ______charged particles (neutrons).

1. James Chadwick (1932)

1. He was an English Physicist.
2. He proved that the nucleus is definitely composed of protons and neutrons through his experiments with ______.

Content Outline: Basic Atomic Structure and Mass (2.2)

1. Atom

1. The smallest particle of an element that stillretains the ______properties of that element.
2. Atoms are composed of 3 sub-atomic particles:

1. Electrons (Thomson proposed)

1. Electrons possess ______electrical charges.
2. Electrons are found orbiting the nucleus of an atom, in what is referred to as the ______ (They move at the speed of light and “create” a cloud-like appearance.)
3. Electrons are 1/1837th the mass of a single proton or single neutron.

1. Protons (Rutherford, Geiger, and Marsden proposed)

1. Protons possess a ______electrical charge.
2. Protons are found clumped together within the ______of an atom.
3. Each proton has a mass of 1 atomic mass unit (AMU) or 1 Dalton (Named after John Dalton.)

1. Neutrons (Rutherford proposed)

1. Neutrons possess ______electrical charged and are therefore referred to as neutral.
2. Neutrons are also found clumped together within the ______of an atom.
3. Each neutron has a mass of 1 AMU or 1 Dalton.

1. Nuclear Forces

1.These are short-range proton-to-neutron OR proton-to-proton OR neutron-to-neutron attractiveforces that help hold together the nucleus of an atom.

2.These forces are greater than the repulsive same chargeelectrical forces exhibit by protons.

1. Atomic Radii

1. This term refers to the relative size of an individual atom of an element.
2. It is measured from the ______to the outermost electron cloud.
3. It is measured in picometers (pm)

1. A picometer is 1.0 x 10-12 meters (so it is very, very small)

1. Charles-Augustin de Coulomb (1785)

1. He proposed ______– attractions that exist between oppositely electrically charged particles (protons & electrons) within a single atom.
2. The forces directly affect the atomic radii of an atom.

1. More ______than electrons = radii shrinking (getting smaller) because the positive charge is greater than the smaller negative charges and pulls them in toward the nucleus.
2. More______than protons = radii increases (getting larger) because the electrons are farther away from the positive nucleus.
3. The Natural state of atoms has protons = electrons; so atomic radii are stable (not changing) for each element.

1. Atomic radii can have an effect on the chemical properties of an element.

1. Atomic Number

1. This term refers to the ______found within the nucleus of an atom for that element.
2. Each element has a ______and identifying number of protons.
3. The atomic number for each element led to the creation of the Periodic Table

1. The Periodic Table was originally created by ______in 1869.

1. He was a Russian Chemist.

1. The atomic number is usually written as superscript (above) the Elements Chemical symbol.

1. Some of the symbols use the Latin term, instead of the English word like Iron, its symbol is Fe for “Ferrum”.

1. Latin is used because it is a “dead” language (will not change over time) and was the original language of science.

1. The Periodic Table was created based upon increasing Atomic Number.

1. Atomic Mass Units (AMU)

1. Also known as the ______
2. This term refers to the total ______of an atom of that element.
3. It is found by ______the number of protons and neutrons together.

1. Each proton OR each neutron has a mass of 1 AMU or 1 Dalton.
2. The Electrons’ mass is insignificant as they are so small (1/1837th that of protons/neutrons).

1. The Atomic Mass is usually written as a ______- (below) the Element symbol.
2. This was based on Carbon-12 as the standard element of measure. It has 12.0 AMU.

1. Isotopes

1. This term refers to atoms of an element that have different masses (AMUs) because they have different numbers of neutrons within the atom; even though it is the same element because they have the same number of protons. (Remember, protons identify the element.)

1. The isotopes behave relatively the same as the natural atom in terms of chemical properties.
2. Some Isotopes are ______(the nucleus is “breaking apart”).

1. To find the number of neutrons:

Start with AMU, subtract the # of protons (atomic number), and that leaves the number neutrons

______= # neutrons

1. How Isotopes are written chemically:

1. Hyphen notation – symbol- number, For example: Carbon-14 OR C-14.
2. Nuclear notation – AMU over Atomic Number symbol, for example 146 C.

1. Nuclide

1. This term is used to refer to the nucleus only (no e- cloud) of an Isotope.

1. Average Atomic Mass

1. As some elements have several isotopes also present in nature, their masses must also be considered to find the average mass for an element (as seen on the Periodic Table).
2. How to calculate the average Atomic Mass of an element:

Step 1: Multiply the AMU for a single isotope by the % found in nature.

Cu 63 – AMU of 63 = 62.93 AMU; so 62.93 x 69.15% (nature) = 62.93 x .6915 = 43.52 AMU

Cu 65 – AMU of 65 = 64.93AMU; so 64.93 x 30.85% (nature) = 64.93 x .3085 = 20.03 AMU

Step 2: Add the all AMUs together.

43.52 + 20.03 = 63.55 AMU

Step 3: Round to two places after the decimal for each isotope calculation in Step 1.

Your Turn: Find the Average Atomic Mass of the following:

1. What is average atomic mass of Lithium if 7.42% exists as 6Li (6.015 g/mol) and 92.58% exists as 7Li (7.016 g/mol)?

1. Magnesium has three naturally occuring isotopes. 78.70% of Magnesium atoms exist as Magnesium-24 (23.9850 g/mol), 10.03% exist as Magnesium-25 (24.9858 g/mol) and 11.17% exist as Magnesium-26 (25.9826 g/mol). What is the average atomic mass of Magnesium?

1. Neon has two major isotopes, Neon-20 and Neon-22. Out of every 250 neon atoms, 225 will be Neon-20 (19.992 g/mol), and 25 will be Neon-22 (21.991 g/mol). What is the average atomic mass of Neon?

Unit 2: Structure of Matter

Content Outline: Atomic Mass and the Mole Concept (2.3)

1. Mole (mol OR n)

1. This is a SI (Le Système International d’ Unités) unit(remember from Unit 1) that is used to represent the amount of a substance.
2. It can be written to show the number of atoms or molecules in a working sample of some element, compound, or molecule, such as sucrose (table sugar) – C6H12O6

1. This concept is very important because scientists, teachers and students cannot work with individual atoms or molecules because they are very, very small and can’t be handled one at a time.
2. So the mole was conceived to represent a ______ of a substance.

1. How to calculate a mole:

1. Determine the total atomic or ______mass of the substance you are working with, using the chemical formula and Periodic Table. (Remember, how to find the Atomic Mass? – Hint…subscript.)
2. Then weigh out, using an electronic balance and weigh boat, that calculated amount.
3. Congratulations, you have just weighed out 1 mole of that substance!

For example: Atomic Mass

Aluminum has an atomic mass of 26.98 AMUs. So you would weigh out 26.98

grams of Aluminum to get a workable amount called a mole.

Molecular Mass Problem

How much does 1 mole of Salt (NaCl) weigh?

The unit for of measurement is g/mol.

1. Molarity

1. Take your 1 mole of a substance and dissolve it in a small amount of distilled water, if it will dissolve, inside a volumetric flask. Then add distilled water to bring the volume to ______of distilled water. You now have a 1 Molar (1 M)solution.
2. Molarity is used for ______solutions.

1. Amedeo Avogrado (1811)

1. He was an Italian physicist.
2. He proposed that the ______ of a gas (at a given temperature & pressure) is proportional to the ______, regardless of the type of gaseous substance used.

1. This eventually was modified to state: That in 1 mole of a substance there will always be ______atoms or molecules present. (That is a massive amount!)
2. This number became known as ______when the French physicist Jean Perrin confirmed and proposed this in 1909 in honor of Avogrado’s work.

1. Jean Perrin would win the Nobel Prize for his work in 1926.
2. The Nobel Prize is Sciences’ highest Award, The Super Bowl trophy in Pro Football.

1. Perhaps your class will celebrate Mole Day on October (10th month) 23 at 6:02 am.

1. Basic Measurements or Unit Conversions involving the Mole concept:

1. More than a mole: the basic concept is: amount you have/ amount of a mole = # of moles

You have 21.6 grams of Boron (B). How many moles do you have?

You have 77.25 grams of Phosphorus (P). How many moles do you have?

1. Less than a mole: the basic concept is: amount you have/ amount of a mole = # of moles

You have 16.03 grams of Sulfur (S). How many moles do you have?

You have 2.43 grams of Magnesium (Mg). How many moles do you have?

1. Conversions from one unit to another unit involving the mole concept:

The basic concept is: Unit given x unit wanted = Unit wanted

unit given

The given unit cancels out and leaves you with the unit wanted.

1. Moles  Atoms/Molecules

a. You have 2.0 moles of Copper. How many atoms of Copper do you have?

b. You have 0.25 moles of Oxygen. How many atoms of Oxygen do you have?

1. Atoms/Molecules  moles

1. You have 1.806 x 1024 atoms of Zinc (Zn). How many moles of Zinc do you have?

1. You have 5.9 x 1022 atoms of Titanium (Ti) How many moles of Titanium do you have?

1. Grams  Moles

1. You have 54.0 grams of Carbon (C). How many moles of Carbon do you have?

1. You have 10.0 grams of Nickel (Ni). How many moles of Nickel do you have?

1. Moles  Grams

1. You have 8.5 moles of Fluorine (F) gas. How grams of Fluorine do you have?

1. You have 0.45 moles of Scandium (Sc) gas. How grams of Scandium do you have?

Unit 2: Structure of Matter

Content Outline: Working Complex Mole Concept Problems (2.4)

1. In most chemistry equations, students need to be able to perform numerous conversions, often referred to as ______.

1. Complex Mole Concept problems.

1. The underlying fundamental concept map (often shown as a mole map)

Atoms/particles  moles  grams

We can go from one end of the map to the other end remembering:

Unit given X Unit Wanted 1 x Unit Wanted 2 = Unit Wanted 2

Unit Given Unit Wanted 1

1. Atoms/Particles  grams

a. You have 9.45 x 1023 atoms of Cesium (Cs). How many grams of Cesium do you have?

b. You have 1.25 x 1023 atoms of Tin (Sn). How many grams of Tin do you have?

1. Grams  Atoms/Particles

1. You have 156.90 grams of Gallium (Ga). How many atoms do you have?

1. You have 380.00 grams of Glucose (C6H12O6). How many molecules do you have?

Unit 2: Structure of Matter

Content Outline: Bohr Model of atoms and Electron Energy (2.5)

1. Niels Bohr (1913)

1. Danish Physicist
2. Proposed the Bohr Model of Atom structure

1. Electrons travel in set paths around the nucleus called ______or ______.
2. Each orbit corresponds with an energy level.

1. Electrons have a natural tendency to occupy the lowest (most stable) energy level first.

1. The lowest energy level (______) is the closest to the nucleus.
2. This is related to ______-- – opposite electrical forces attract.
3. The farther away from the nucleus the greater the Potential Energy for that electron.
4. The closer to the nucleus the less Potential Energy for that electron.

1. Electrons can absorb energy (absorption) from their surroundings from another energy source, such as sunlight energy (A.K.A electromagnetic energy).

1. Electrons that gain energy (absorption) are said to be “excited”.
2. Electrons that lose energy (______)emit light as they return to a more stable (less energy) grounded state.
3. The unit of light energy is referred to as a ______.
4. The unit of measurement for the energy lost OR gained by an atom is a ______.

1. Electrons as ParticlesWaves

1. Electrons can move as particles around the nucleus because they have mass (if ever so small).
2. Electrons, as they are moving, move in wave-like fashion (like waves in the Gulf of Mexico… up, down, up, down)
3. Properties of waves

1. Wavelength (λ)

1. This is defined as the distance between ______(such as crest –top or ebb- bottom) on adjacent waves.

1. As it is distance, some unit of measurement of distance (meter) is used for wavelength, usually nanometer (nm OR 10-9).

1. Frequency (v)

1. This is defined as the ______ that pass a given point in a specified time, usually seconds.
2. Frequency is expressed in Hertz (Hz) or waves/sec.

1. Heinrich Hertz defined 1 wave/sec = 1 Hertz.

1. Speedof light (c)

1. Electrons travel at the speed of light.
2. Waves are measured against the speed of light (electromagnetic radiation).
3. C = ______ is the equation for the speed of light.

1. As light speed never changes, it is considered to be a constant at 3.00 x 108 m/sec.
2. The properties of light are inversely proportional.

α. As wavelength decreases, frequency increases.

b. As wavelength increases, frequency decreases.

d. As electrons gain more energy, they travel faster and get farther from the nucleus.

1. Electromagnetic Spectrum (Light Energy)

1. This term refers to the whole spectrum (variations) of ______.

1. ______ is used to define the wave-like movement of light particles.
2. Light moves at 3.00 x 108 m/sec.
3. The electromagnetic spectrum includes: x-rays, microwaves, visible (white) light, ultra-violet light, infrared light, and radio waves.

Unit 2: Structure of Matter

Content Outline: Photoelectric Effect and Emission Spectrum (2.6)

1. Photon

1. Albert Einstein proposed this concept in 1905.
2. A Photon is a unit of light energy having no mass and possessing a single quantum of energy.

1. Quantum

1. Minimum quantity of energy that can be lost or gained by an atom.
2. Proposed by German physicist Max Planck in 1900.
3. This sets the field of Quantum Physics (nanoscale physics) in motion.
4. Planck wins the Nobel Prize in 1918 for this work.

1. Photoelectric Effect

1. The ______- (ejection) of an electron from a metal surface when light shines on the surface.

1. This shows a direct connection between light and light possessing energy.
2. This light energy is Einstein’s photon.

1. The light has to be of a ______ in order for the effect to take place.
2. Each metal requires a ______frequency of light.

1. Planck’s Constant Theory

1. This tries to explain the Photoelectric Effect by proposing a relationship between a ______of energy and the ______of radiation. Remember, light is considered electromagnetic radiation, so the frequency changes with the various forms of light/radiation. This is called the ______.

1. E = hv

1. E = energy for a quantum of radiation.

1. It is measured in Joules (J).

1. v = frequency of the radiation.

1. measured in waves/sec. (s) -1 or Hertz

1. h = Planck’s Constant

1. Defined as 6.626 x 10-34 J * s (* = times)

1. Planck-Einstein Relation

1. Albert Einstein expanded on Planck’s work in 1905.
2. He proposed that light has a combination of wave properties and particle properties.

1. Each particle of light carries 1 quantum of energy

1. E = hv (Planck’s version) then becomes Ephoton = hv (Einstein’s version).
2. Matter can only absorb Electromagnetic radiation (light) in whole number (1, 2, and so forth) quantities of photons.
3. In order for a single electron to be emitted from the metal surface, the electron must be struck by a ______- possessing at least the minimum amount of energy to eject the electron.

1. This minimum amount of energy is directly related to the ______

1. The ______ the frequency – the more possible to emit an electron from the surface.
2. The ______the frequency – the less likely to emit an electron from the surface.

1. Different metals require different frequencies for the Photoelectric effect to take place.

1. Emission Spectrum of Hydrogen

1. This is an expansion of the ______Tube experiment.

1. It uses Hydrogen gas (which glows pink) and a glass prism (triangular shaped piece) placed in the path of the light ray.

1. The light ray split into 4 different colors (red, green, blue, and purple).

1. These become known as the ______spectrum.
2. Each color represents a fixed quantity of energy for an excited electron.
3. It is later added for other frequencies of light, such Infrared and Ultra-violet.

1. These experiment set the groundwork for the Modern Quantum Atomic Model of atoms.

Unit 2: Structure of Matter