Naming Compounds!

Today during class, Mr.Doktor taught us how to name compounds and we went over the following:

Chemical Nomenclature

• Today the most common system IUPAC- most chemicals.
- Ions
- Binary Ionic
- Polyatomic Ions
- Hydrates
- Molecular Compounds
- Acids/Bases

Chemical Formulas

• Be aware of the difference between ion and compound formulas.

Examples

Zn^2+…………. ion charge (Zinc Ion)
BaCl2…………..number of ions (Barium Chloride)

Non-Metal Ions
F-………………Fluoride Ion
N^3-……………Nitride Ion
O^2-……………Oxide Ion

AlF3………….Aluminum Fluoride
Na2O…………Sodium Oxide
Fe2S3…………Iron (III) Sulfide

Multivalent Ions

• Some elements can form more than one ion.
- Iron Fe^3+ or Fe^2+
- CopperCu^2+ or Cu^1+
• The top number on the Periodic Table is more common.
• IUPAC uses roman numerals in parenthesis to show the charge.
• Classical(i.e. old) systems uses Latin names of elements and the suffixes –ic (larger charge) and –ous(smaller charge)

Examples

Fe^3+…..Ferric Oxide (Fe2O3)
Fe^2+….Ferrous Oxide (FeO)

Other Classical Names

• Ferr-Iron
• Cupp-Copper
• Mercur-Mercury
• Stann-Tin
• Aunn-Gold
• Plumb-Lead
• Wolf-Tungsten
• Argent-Silver

Old Greek Way

Argentous OxideAg2O
Cuppric Sulfide CuS

IUPAC

CuCl2- Copper (II) Chloride
Cr2O3- Chromium (III) Oxide

Hydrates

• Some compounds can form lattices that bond to water molecules.
• Copper Sulfate
• Sodium Sulfate
• These crystals contain water inside them which can be released by heating,

• To Name Hydrates:

- Write the name of the chemical formula.
- Add a prefix indicating the number of water molecules (mono=1, di=2, tri=3 etc.)
- Add hydrate after the prefix.

Examples

Cu (SO4) •5H2O……….Copper (II) Sulphate Pentahydrate (5 Water)
Li (ClO4) •3H2O………Lithium Hypochlorite Trihydrate (3 Water)

Electron structure (Electron dot diagrams)

In today class we learned how the electrons are present in each element in either basic element form or as an ion. In an ion the element is either positively or negatively charged and either lose or gain an electron depending on the charge of the element. A negatively charged ion would gain electrons and a positively charged ion would lose electrons.

Drawing Electron Dot diagrams:
• To represent the nucleus, write the atomic symbol
• For individual elements determine the number of valence electrons
• Electrons are represented by dots around the symbol
• Four orbitals (one of each side of the nucleus) each holding a max of 2 electrons
*Each orbital gets 1 electron before they pair up
e.g.

Lewis diagrams for Compounds and Ions:
• In covalent compound, electrons are shared
1. Determine the number of valence electron for each atom in the molecule
2. Place atoms so that valence electrons are shared to fill each orbital
e.g.

Ionic Compounds:
• In ionic compounds electrons transfer from one element to another
• Determine the number of valences on the cation is more than to the anion
• Draw [] around the metal & non metals
• Write the charges outside the brackets
e.g.

Trends on the Periodic Table

Elements close to eachother on the periodic table display similar characteristics

There are 7 Periodic trends: Reactivity, Ion Charge, Melting Point, Atomic Radius, Ionization energy, Electronegativity, & Density.

Reactivity
Metals and non-metals show different trends.
- The most reactive metal is Francium
- The most reactive non-metal is Fluorine

Ion Charge
Elements ion charges depend on their group/column.

Melting Point
Elements in the center of the table have the highest melting point
- Noble gases have the lowest melting points
- Starting from the left and moving right, melting point increases (until the middle of the table.

Atomic Radius
Radius decrease to the up and the right
- Helium is the smallest atomic radius
- Francium is the largest atomic radius

Ionization Energy
The energy neededx to completely remove an electron from an atom
- it increases going up and to the right
- all Noble Gases have high ionization energy
- Helium (non-metal) has the highest ionoization energy
- Francium (metal) has the lowest ionization energy
*OPPOSITE trend from Atomic Radius

Electronegativity
Refers to how much atoms want to gain electrons
*SAME trend as Ionization Energy
- Francium: wants to gain electrons

Density
How compact an atom is.

-From top to bottom, density increases

Isotopes and Atoms!

Today we learned about atoms and isotopes. No atom is perfectly the same; some atoms have variations of mass (due to the difference in neutron number) and these are called isotopes. We also learned how a spectrometer works and how it gets the average of the different isotopes.

Important Definitions
Atomic Number - Number of protons in an atom
Isotopes - Same atomic number but different mass
Spectrometer - are used to determine the abundance and mass of the isotopes of elements

Isotopes of Hydrogen


Spectrometer: Calculating averages


We have in the image shown above the relative abundances of europium. To find the average mass we need to do a series of calculations...

Step 1: Multiply the percentage (convert to decimals) with the mass number
151(0.478) + 153(0.522) = Average Mass
Step 2: Add the results
72.178 + 79.866 = Average Mass
Step 3: You get your average
152.0 = Average Mass of Europium

*Check the periodic table! Europium has an atomic mass of 152.0!

Quantum Mechanics

Today in chemistry class we compared the bohr theory and the quantum theory. Mr. Doktor then taught us about electron configuration and the different orbitals. It was confusing but after doing a lot of examples it was clear and we understood what the it was about.


Bohr Theory
- The electron is a particle that must be in orbital in the atom.

Quantum Theory
- The electron is a cloud of negative charge or a wave function.
- Orbitals are areas in 3D space where the electrons most probably are.
- The energy of the electrons is in its vibrational modes- like notes on a guitar string.
- Photons are produced when high energy modes change to lower energy modes.

S Orbitals

- Each orbital holds 2 electrons

P Orbitals

- There are 3 suborbitals
- Each contains 2 electrons
- Total electrons = 6

D Orbitals

- There are 5 suborbitals
- Each contains 2 electrons
- Total electrons = 10

F Orbitals

- There are 7 suborbitals
- Each contains 2 electrons
- Total electrons = 14

Bohr Diagrams

In today's class, it was just a continuation of last class and just went into further study of what elements the main Bohr diagrams are made up of and how to draw them. The drawing of a bohr diagram was just a review from grade 9 so it was pretty easy.

In a bohr diagram:
• PROTONS=ATOMIC NUMBER
• NEUTRONS=ATOMIC MASS
• Atoms are electrically neutral

There are two different models for fluorine that can be used to describe electron configuration:
1. Energy Level Model:

2. Bohr Model:

Bohr Model

Todays class we had a lesson to learn more about the Bohr Model & deepen our understanding of Niels Bohr's theory.

Bohr's Theory
- Electrons exist in orbitals
- When they absorb energy they move to a higher orbital
- As they fall from a higher orbital to a lower one they release energy as a photon of light.

BOHR (1920s)
- Rutherfords model was inherently unstable
protons & electrons should attract eachother
- Matter emits light when it is heated
- Light travels as photons and the energy photons carry depends on their wavelength
seperating white light with a diffraction of wavelengths

Each line represents a photon of light emitted from the excited atom

- Bohr based his model on the energy emitted by different atoms
- Each atom has a specific spectrum of light & each spectrum represents an element
Bohr suggested that electrons occupy shells or orbitals to explain this emission.

We also watched a video on YouTube called the Double Slit experiment by Dr Quantam, which shows particles that go through a slit and form a pattern on a backboard, and then figuring out why there is an interference pattern on the back wall when there are 2 slits made to form a pattern on the backboard.

Atomic Thoeries!

Today we learned about many atomic thoeries throughout history. An atomic thoery relates idea/thoeries about small particles such as atoms, protons, and electrons.

These are some of the atomic thoeries we learned today:

Atomic Thoery	Main Features	Diagram	Shortcomings/Problems
Democritus 300 BC	Talks about the atom as the smallest particle of matter. Defines the atom as an indivisible particle Explains certain natural occurrences such as the existence of elements	Atom the indivisible particle Atomos (in ancient Greek) means "that which cannot be further broken down into smaller pieces".	Does not give a scientific view of the atom only a conceptual definition Does not talk about subatomic particles (Electrons, Protons, Neutrons)
John Dalton 1800s	Explains a lot of chemical properties such as how atoms combine to form molecules Explains chemical change better than the Particle Theory Confirms the basic Laws of Chemistry: Conservation of Mass & definite Proportions	The solid sphere model Atoms are seen as solid, indestructible spheres (like billiard balls)	Does not include the existence of the nucleus Does not explain the existence of ions or isotopes Does not talk about subatomic particles (Electrons, Protons, Neutrons)
J.J. Thompson 1850s	Infers on the existence of electrons and protons Introduces the concept of the nucleus Infers on the relative nuclear density and atom mass of different atoms	The raisin bun Model or the chocolate chip cookie model : Atoms are solid spheres made-up of a solid positive mass (or core) with tiny negative particles embedded in the positive core.	Does not explain the existence of electrons outside the nucleus does not explain the role of electrons in bonding Does not talk about neutrons therefore can't explain radioactivity and the existence of isotopes
Rutherford 1905	First real modern view of the atom Explains why the electron spins around the nucleus Proposes that the atom is really mostly empty space	The Planetary Model Famous Gold Leaf Experiment proves that the nucleus is positive and the electrons are outside the nucleus.	Does not place electrons in definite energy levels around the nucleus Doesn't include neutrons in the nucleus Does Not relate the valence electrons atomic charge
Neils Bohr	Explains the role of valence electrons in bonding Relegates the number of valence electrons to the Periods of a periodic table Fully explains ionic and covalent bonding Places electrons in definite energy levels 2 e- in the first 8 e- in the second 8 e- in the third	Electrons in Definite energy Levels around the nucleus Used atomic spectra to prove that electrons are placed in definite orbitals (called shells) around the nucleus.	It does not explain the shapes of molecules or other abnormalities that result form unevenly shared pairs of electrons (such as the abnormal behaviour of water, the difference in Carbon-Carbon Bonds between diamond and graphite etc..)

Density & Graphing!

Today in Chemistry Class, we learnt about Density and Graphing. We learnt that the density of an object is its mass divided by its volume or the equation of:

*Density is usually expressed in Kg/L, Kg/m3, or g/cm3

Graphing (5 important things to remember)

1. Labeled axis

2. Appropriate scale

3. Title

4. Date Points

5. Line of Best Fit

3 things can be done while we work with graphs and that is

* Read the graphs

* Find the slope(rise/run)

* Find the area of the graph

example of a properly constructed graph: