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SCH 4U Name_______________________

CHEMISTRY REVIEW Study Guide: Vocabulary

€ Matter, pure substance, mixture, compound, element, homogenous mixture, heterogeneous mixture

€ metal, non-metal, alkali metals, alkaline earth metals, transition metals, halogens, noble gases, metalloids

€ atom, proton, neutron, electron, molecule, valence

€ physical property, physical change, chemical property, chemical change

€ reactants, products

€ Law of conservation of Mass/Energy

€ synthesis, decomposition, single displacement, double displacement, combustion, neutralization

€ acid, base, neutral, strong, weak, indicators

€ solubility, soluble, low solubility, insoluble

Compare and Contrast, Label

€ pure substance vs. mixture

€ homogeneous vs. heterogeneous

€ properties of metals, non-metals, metalloids

€ properties of protons, neutrons, and electrons

€ properties of ionic and molecular compounds

€ intermolecular forces

€ physical vs. chemical property

€ physical vs. chemical change

€ qualitative vs. quantitative observation

€ properties of acids, bases, and neutral substances

€ strong vs. weak acids, strong vs. weak bases

€ radiation diagrams

€ properties of light, electricity, electromagnetic radiation, energy

Diagrams, Graphs, Charts, and Lists

€ Matter/pure substance/mixture chart

€ Bohr Rutherford Diagrams

€ Lewis Dot Diagrams

€ VSEPR shapes of molecules

€ indicator colors: red litmus, blue litmus, phenolphthalien, bromothymol blue

€ list of polyvalent metals

€ properties of gases

€ activity series

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Calculations and Procedures

€ calculate the number of protons, neutrons, and electrons in an element, given atomic number, mass number (and vice versa)

€ name ionic and molecular compounds

€ write with formulas for ionic and molecular compounds

€ write word equations and chemical equations

€ balance a reaction equation

€ count the number of atoms in a molecule from its molecular formula

€ stoichiometry & limiting reactants

€ average atomic mass

€ concentration calculations

€ acid/base calculations

€ % yield

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UNIT 1 - STRUCTURE AND PROPERTIES OF MATTER

CHAPTER 3 - ATOMIC THEORY REVIEW (PART 1- Matter and The Atom) Chemistry - the study of matter, its composition, structure, properties, and changes

Type of Matter Definition Example Pure Substance Mixture Homogeneous Mixture Heterogeneous Mixture Element Compound

Important Chemistry Laws and Theories Particle Theory of Matter The Particle Theory of Matter states that: 1. Matter is made up of tiny particles (atoms & compounds ) 2. Particles of Matter are in constant motion. 3. Particles of Matter are held together by very strong electric forces 4. There are empty spaces between the particles of matter that are very large compared to the particles themselves. 5. Each substance has unique particles that are different from the particles of other substances 6. Temperature affects the speed of the particles. The higher the temperature, the faster the speed of the particles. Law of Conservation of Mass (Lavoisier) The law of conservation of mass states that mass in an isolated system is neither created nor destroyed by chemical reactions or physical transformations. According to the law of conservation of mass, the mass of the products in a chemical reaction must equal the mass of the reactants. Law of Definite Proportions or Law of Constant Composition (Proust) A chemical compound always contains exactly the same proportion of elements by mass.

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The Atom & Subatomic Particles (protons, neutrons, electrons)

Mass Number - neutrons + protons in the nucleus of an atom Atomic Number - number of protons in one atom of an element Atomic Mass: a weighted average atomic mass of all the natural isotopes of an element. Isotopes: same atomic number but a different mass number.

� Changing the number of protons produces a different element . � Changing the number of electrons produces an ion . � Changing the number of neutrons in an atom produces a different isotope of the

element. Radioisotopes - isotopes of an atom which emit radiation. Occur because the nucleus is unstable due to neutron-to-proton imbalance (α and β decay) or excess energy (γ decay).

Types of Radiation (Radioactive Decay):

o alpha - emits alpha particles (helium nuclei); particles can be stopped by paper

o beta - emits beta positive or beta negative particles (same size as an electron); particles may be stopped by aluminum

o gamma - electromagnetic radiation; highest energy; stopped by lead

Example: hydrogen-2(deuterium) and hydrogen-3(tritium) are radioactive isotopes of hydrogen

mass charge location

proton (p+) neutron (n0)

electron (e-)

Low neutron to proton ratio High neutron to proton ratio Excess energy

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The Periodic Table: The periodic table is divided into two major groups: metals and nonmetals. Metalloids have properties of both metals and non-metals and can be found on either side of the staircase that starts between boron and aluminum. To simplify, everything to the left of the staircase is considered a metal, and everything to the right is considered a non-metal.

Elements are arranged in rows called periods, which represent energy levels in an atom. Elements are also arranged by increasing atomic number, which represents the number of protons in an element's nucleus. This number also represents the number of electrons in a neutral atom.

a) representative elements These groups are divided into families (vertical columns) and these elements have similar physical and chemical properties. This is because the elements in these families have the same number of electrons in their outer most shells.

b) transition metals These metals can react to form different ions because of differences in their valence shell structure.

Other information given on the periodic table Average Atomic Mass - a weighted average of all the isotopes of an element by naturally occuring percentages on earth. Electronegativity - the tendency of an atom to attract a bonding pair of electrons towards its nucleus Common ion charges - the charges each element takes when forming ions or bonding.

Properties of Metals Properties of Non-metals Properties of Metalloids, - solid at room temperature

- shiny (mostly) - brittle (mostly) - semi-conductive

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Atomic or Isotopic Notation mass number element atomic number symbol EXAMPLE Helium-4 Helium-5 (heavier isotope)

4 5 He He 2 2 Chemically, helium4 would react exactly the same as helium-5. Ionic Notation: EXAMPLE mass number +/- charge Sodium-23

23 Na1+ 11 element

Practice Questions Complete the following chart.

Atomic or Isotopic Notation

Name Proton# Electron # Neutron # (mass#-proton#)

a)

Calcium-42

b)

13

14

c)

Zinc-63

28

33

d) 127 I -

53

e)

Oxygen-16

10

f) 56 Fe2+ 26

proton # = 2 electron # = 2 neutron # = 4-2=2

proton # = 2 electron # = 2 neutron # = 5-2=3

proton # = 11 electron # = 10 (lost 1 e-) neutron # = 23-11=12

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Representing Atoms Bohr-Rutherford Diagrams show all the electrons around an atom's nucleus, arranged in concentric shells. The inner shell can hold 2 electrons, and the outer shells can hold eight. Steps: 1. Draw a circle to represent the nucleus. Inside write the number of protons and the number of neutrons in the nucleus. 2. Draw the first electron shell around the nucleus. This holds two electrons (place together or opposite each other). 3. Draw the next electron shell outside of the first. Subsequent electron shells are filled by placing one electron on each of four sides clockwise, and then doubling up (see diagram). Example: a complete Bohr Rutherford diagram for potassium Draw Bohr Rutherford Diagrams for the following elements

Lithium Fluorine Aluminum

Nitrogen Argon Carbon

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Lewis Dot Diagrams show the element's chemical symbol and number of valence electrons. Valence Electron - an electron in the outmost energy level or orbit. These electrons are important to focus on when we are making bonds and forming ions. Octet Rule - Atoms will combine with other atoms so that there will be a transfer or sharing of electrons so that all the atoms involved will have a stable octet (8 valence electrons). This can also be achieved through ion creation. Drawing Lewis Dot Diagrams (aka Electron Dot Diagrams) 1. Chemical symbol goes in the middle 2. Electrons go around the 4 sides (one per side) and then double up. Order of sides does not matter. Examples

Try -> Drawing Lewis Structures for Ions 1. All Lewis dot diagrams for ions are surrounded by a set of square brackets [ ]. The charge of the ion is written as a superscript outside of the brackets. 2. Metal atoms lose electrons to form positive cations. Their valence shell becomes empty. Charge number = # e- removed 3. Non-metal atoms gain electrons to form negative anions. Their valence shell becomes full. Charge number = # e- gained. Examples

Draw Lewis Dot Diagrams for the following ions

sodium ion

nitrogen ion

carbon ion

potassium ion

aluminum ion

- end of review -

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3.1-3.3 The History of Atomic Theory Timeline of Contributions to Atomic Theory

Leucippus and Democritus - philosophers - developed a theory that if you kept cutting bread, you would eventually reach an indivisible particle - named the atom "Atomos" (a = not, tomos= divisible)

Antoine Lavoisier

- did experiments on the masses of reactants and products - Law of Conservation of Mass: the masses of the reactants of a chemical reaction equal the mass of the reactants

John Dalton

- Billiard Ball model of atom - Atomic Theory: 1. all matter is made of atoms 2. atoms of the same element have the same properties 3. compounds are formed by a combination of different atoms 4. chemical reactions rearrange atoms

Dimitri Mendeleev

- sorted different elements by their properties - published the first periodic table

William Crookes Eugen Goldstein

- Crookes invented the cathode ray tube (two metal plates in a vacuum tube with electricity run through them) - Crookes discovered a negative ray from the anode - Goldstein discovered a positive ray from the cathode

Heinrich Hertz

- discovered that you could make sparks by shining a UV light at a metal - called the "Photoelectric Effect"

1887 1875 1886

1869 1803 1785 5th cent.

BCE

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Timeline of Contributions to Atomic Theory

Max Planck

- studied black body radiation - thermal radiation has a different color in a continuous spectrum depending on its temperature - coined the term quantum

J.J. Thompson - used a cathode ray tube - discovered the electron (areas of the atom are negatively charged) - developed the plum pudding/ blueberry muffin/ chocolate chip cookie atomic model

Marie and Pierre Curie - discovered radioactive decay in elements - discovered radium and polonium - coined the term "radioactivity"

Robert Millikan - Millikan oil drop experiment - discovered the charge of a single electron

Albert Einstein - explained the photoelectric effect - described light as containing discrete packets of energy (quanta) rather than continuous waves - light is both a wave and a particle

Ernest Rutherford - Rutherford gold foil experiment - showed that there was great space within the atom, but that the atom had a solid, positive core - discovered the proton - planetary atomic model

1911 1909 1905 1898 1897 1890

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Timeline of Contributions to Atomic Theory

Neils Bohr

- studied the emissions spectra of hydrogen gas - found it contained very specific colors of light, and related it to the energy levels of electrons - created the solar system model of the atom

Frederick Soddy

Francis Aston

- both are credited with the discovery of isotopes - Soddy discovered that radioactive isotopes had different masses - Aston used the mass spectrometer to discover more

Louis de Broglie

- double slit experiment - proved that light behaved both like a particle and a wave - contributed to quantum theory

Werner Heisenberg

- created quantum mechanics - came up with the Heisenberg uncertainty principle (the speed and position of an electron cannot both be known at once)

Erwin Shrödinger

- developed the electron cloud theory of the atom - came up with the Shrodinger equation, a wave equation which relates the position and energy of electrons

James Chadwick

- discovered the neutron by bombarding beryllium with alpha particles

1932 1926 1925 1924 1913 1919

1913

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HISTORY OF ATOMIC THEORY NAME, OCCUPATION & EXPERIMENT THEORIES OR DISCOVERIES MODEL OF THE ATOM Democritus (Philosopher) (400 BC) - Experiments: none (he was a philosopher)

Theory of the Atom: when matter is divided into smaller and smaller pieces, a finite limit known as the atom is ultimately reached.

- none

John Dalton (Scientist) (1803) - Experiments: lots of experiments with combining atoms of different elements Observations:

• Elements cannot be broken down • Compounds can be broken down into elements • Law of Conservation of Mass

Dalton's Atomic Theory: 1. All matter is made of tiny indivisible particles called atoms. (FALSE: atoms have subatomic particles). Atoms cannot be created or destroyed. (FALSE: radiation can destroy atoms) 3. All atoms of a particular element are identical. (FALSE: atoms have isotopes with different masses) 4. Compounds are formed through the combination of elements. STILL TRUE 5. Chemical reactions occur when atoms in compounds join together or separate to form new compounds. STILL TRUE

Billiard Ball Model

JJ Thompson (Scientist) (1896-97) Cathode Ray Tube

Discovery: Electrons! - The cathode of the cathode ray tube, emits small particles that deflect towards positive plates and away from negative plates. Therefore the particles have -ve charge! - There are particles smaller than atoms that can be emitted from atoms and have a negative charge

Plum Pudding Model

Heinrich Hertz(1887) Max Planck(1900) Albert Einstein(1905)

Discovery: Quantum Energy! -Heinrich Hertz discovered that when you shine light on a metal, light (electrons) are emitted, however, only higher frequencies of light produced this effect - Max Planck discovered black body radiation. He heated solids until they glowed and measured the frequency of energy given off. He found that the intensity of the light increased with heat, but the frequency (wavelengths) did not. They were restricted to infrared through ultraviolet. E=nhf (Matter can only release energy in whole number multiples of certain frequencies). - Einstein later explained that this indicated that light energy is carried in discrete packets called quanta.

Ernest Rutherford (Scientist) (1911)

Results: 99% go right through (lots of space) 1% deflect/bounce

Discovery: The Nucleus! - charges and matter are not uniformly distributed throughout the atom - there is a small, dense, positive core (the nucleus), surrounded by a lot of empty space

Planetary Model

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NAME, OCCUPATION & EXPERIMENT THEORIES OR DISCOVERIES MODEL OF THE ATOM Neils Bohr (Scientist) (1913)

When excited, hydrogen gas in a discharge tube (like a neon light), emits light. when that light passes through a prism, each color bends a different amount, separating light of different wavelengths (colours). Hydrogen was found to emit only 4 specific wavelengths.

Discovery: Energy Levels!

Bohr's atomic model stated that: - Electrons only assume certain orbits around the nucleus - Energy is associated with each orbit, exactly the same amount required to make the same jump from one orbit to another, a quantum - Light is emitted when an electron jumps from a higher orbit back down to its original lower orbit

Solar System Model

Louis de Broglie (Scientist) (1924)

Discovery: Electron Wave-Particle Duality! De Broglie stated that: - Electrons could act as both particles and waves; his double split experiment provided evidence for this Werner Heisenberg (1927) came up with: - The Uncertainty Principle which stated that it is impossible to determine elementary particles' locations and momentum Erwin Schrödinger's theory (1926) stated that: -Electron density clouds, or orbitals, existed - The more dense the orbitals were, the higher the probability of finding an electron in that area

Electron Cloud Model

waves can cancel each other out or create areas of high intensity when passed through slits. (interference pattern) particles shot through two slits were expected to travel in a straight line making a two-line pattern electrons created an interference pattern like light waves, so electrons are a particle that travels like a wave.

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3.4 Quantum Numbers There are 4 quantum numbers, n (the principle quantum number), l (the orbital quantum number), ml (the magnetic quantum number) and ms (the spin quantum number). Three of these numbers come from Schrödinger's wave equation for the probability of finding an electron within a particular space around the nucleus of the atom. Every electron within an atom will have a unique set of quantum numbers.

n (the principle quantum number) {1,2,3,4,5,6,7...} - the energy level or "shell" that the electron belongs to - if we are finding quantum numbers for the outer (valence) shell of an element, the n number is the same as the number of the period that element is in. (eg, He:1, C:2, Mg:3) l (the angular quantum number) {0,1,2,3,4,....} - the orbital type that the electron belongs to (also called a subshell) - matches neatly with the organization of the periodic table - each orbital has its own distinct three-dimensional shape value of l 0 1 2 3 letter used s p d f name sharp principle diffuse fundamental

Shapes

s block - alkali metals and alkaline earth metals

d block - transition metals

p block - non-metals and metalloids

f block - lanthanides (4f) and actinides (5f)

size, nodes, and energy

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ml (the magnetic quantum number) {...-3,-2,-1,0,1,2,3...} - the number that describes the orientation of the orbital that the electron belongs to in space relative to the other orbitals of the same type in the atom

s orbital

1 orientation

ml = 0

p orbital

3 orientations

ml = -1, 0, 1

d orbital

5 orientations

ml = -2, -1, 0, 1, 2

f orbital

7 orientations

ml = -3, -2, -1, 0, 1, 2, 3

ms (the spin quantum number) {-1/2, +1/2} - proposed by Goudsmit and Uhlenbeck to account for emissions spectra of atoms - atoms have a magnetic property when placed in an external magnetic field - spinning charges produce magnetic moments - atoms within the same period, orbital, and orbital orientation have two spin states

Relating quantum numbers back to the 2n2 electron shell formula

Quantum Number Summary

Practice Questions What are the quantum numbers for the last electron in... a) a carbon atom b) a copper atom c) a bromine atom d) a sodium ion

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Practice Questions (#3-12)

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3.5 Energy Level Diagrams Orbital Diagrams and Electron Configurations Now that we know about the structure of an atom, and quantum numbers, we have two new ways of representing the electrons in an atom, energy level diagrams and electron configurations. Energy Level Diagram a diagram that represents the relative energies of electrons in an atom

Fig 1. Energy Level Diagram for Oxygen Fig 2. Aufbau Diagram

Aufbau Principle: Atoms can be 'built up' by the addition of electrons, which fill orbitals starting at the lowest energy level before filling higher energy levels. You can use an Aufbau diagram or your periodic table to determine the order of orbitals with increasing energy Hund's Rule: Electrons will enter empty orbitals of equal energy, when they are available.

Pauli Exclusion Principle: No two electrons in an atom have the same set of quantum #s.

Orbital Diagrams and Electron Configurations

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Atoms with Unexpected Electron Configurations Chromium Expected: [Ar] 4s23d4 Actual: [Ar] 4s13d5 Copper Expected: [Ar] 4s23d9 Actual: [Ar] 4s13d10 New Rule: To make a full, or half-full d shell, promote 1 electron from the s shell.

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Applications of Electron Configurations Explaining ions of multivalent metals Many of the transition metal elements can form ions with two or more different charges. For example, Pb (lead) can form ions of 2+ and 4+ charges. Neutral lead atom: [Xe] 6s24f145d106p2

Pb 2+ atom: [Xe] 6s24f145d10 (full s,f, and d shells, empty p shell) Pb 4+ atom: [Xe] 6s04f145d10 (full f, and d shells, empty s and p shell) Magnetic Properties of Transition Metals Some transition metals are strongly magnetic, but some are only weakly magnetic. Why? Ferromagnetism - When a metal is exposed to a magnetic field, and this magnetic field permanently induces a magnetic property within a metal. Examples: iron, nickel, cobalt. Why? These substances have unpaired electrons AND are arranged clusters of atoms called domains, where electrons line up in the same direction. Different domains line up in different directions, but when a magnetic field is applied, they all line up in the same direction. Paramagnetism - When a metal is exposed to a magnetic field, it experiences weak attraction due to unpaired electrons, but has no domains. Examples: aluminum Diamagnetism - When a metal is exposed to a magnetic field and is very weakly repelled, due to the absence of unpaired electrons, i.e. all electrons are paired. Example : gold

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Chapter 3 Review - Self Quiz, Page 185 (all)

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Chapter 3 Review, page 186

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Chapter 3 Review, page 187

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Chapter 3 Review, page 187