Prezi Project
An atom is composed of a dense nucleus, which contains protons and neutrons.
- Protons = Positive Charge
- Neutrons = Neutral Charge
- Electrons = Negative Charge
Isotopes = Atoms of the same element but have different atomic mass. They contain varying amount of neutrons.
Periodic Trends
Metals and Metallic Bonds
Conversion
Introduction to Chemical Quantities
Mole: Fundamental counting unit in chemistry.
- 6.02 x 10^23 atoms
- Moles are a given quantity that can be directly compared in the visible level
- Equal to the atomic mass (grams) of any element or molecule.
Interstitial and Substitutional Alloys
The Electron Sea and Delocalized Electrons:
- When metallic atoms bond, they're bonded through mutual sharing of the electron sea.
- Each atom's valence electrons are said to be delocalize, which they flow freely around the nuclei.
Metallic bonds: Collection of metallic ions (cations) living in a group of free-flowing electrons.
- Conductive
- High density
- Malleable
- Ductile
- Magnetism
- Relative softness
- Low phase-change
- 1 min = 1 sec
- 1 hr = 60 min
- 1 day = 24 hrs
- 1 m = 100 cm
- 1 in = 2.54 cm
- 1 ft = 12 in
- 1 ml = 5,280 ft
Percent by Mass
A given reaction will require a certain number of atoms or molecules to reactant.
- Since different elements have different atomic masses, the same mass of two different substances doesn't alwyas mean the same amount of two different substances.
- Alloy: Compound composed of two or more metals f of a metal and another non-metal that exhibits metallic bonds
- Substitutional Alloys: Formed when atoms of one type of metal are replaced with another metal of similar size.
- Can only be formed if replacement element is of similar size to the original.
- Institutional Alloys: Formed when spaces between the atoms of a metal are filled using another metal or element
- Additional element must be sufficiently smaller than the original metal so that it may fit in the gaps created by metals atom.
According to law of constant composition, all molecules of substance have same ratio elements.
Empirical and Molecular Formulas:
- Molecular Formula: Tells exactly what something is made off
- Empirical Formula: What the simplest ratio of elements present for a substance is.
- Multiple substances will have the same empirical or molecular formula, but each substance will always have a different molecular formula.
London Dispersion Forces
Steps
Hydrates
Hydrates: Compounds made up of water molecules trapped within a crystalline lattice
- All hydrate chemical formulas can be written as a simple ratio of one salt; x water
- Possible for a given salt to form more than one type of hydrate.
To calculate the empirical or molecular formula...
1. Assume 100 grams.
- Only assume 100 unless given grams at the start.
2. Calculate how many moles of each substance you have. Write as ratio.
3. Simplify your ratio by dividing all terms by smallest value present.
4, Simplify ratio again (if necessary) by multiplying a constant to get rid of fractions.
5. Use new ratio to write empirical formula.
6. Compare empirical formula to the atomic mass (if given) to identify the molecular formula.
Intermolecular Forces
Intramolecular Bonds: Ties atoms together to form a molecule within.
- Always stronger than intermolecular bonds.
Intermolecular Forces: Attaches molecules to each other between (IMFs).
- Molecules present in a substance must be held together
3 categories of IMFs:
Dipole-Dipole
London Dispersion Forces
Hydrogen Bonds
- Instantaneous attachment; only attaches for a second.
- Created by induced migration of electrons in an atom or molecule.
- Results in slightly negative and slightly positive side of the molecule being created.
- chain reaction
- Weakest of the IMFs
- If one molecule or atom exhibits an LDF, will induce LDF in its neighboring molecules or atoms.
- Strength increases with increased number of electrons.
- Greater the mass, greater the amount of electrons
- Specific type of dipole-dipole
- Must be made from molecules.
- Only exist between molecules that have slightly positively-charged hydrogen and slightly negatively-charged nitrogen, oxygen, or flourine.
- Uncharacteristically stronger than other dipoles.
- Strongestnof the IMFs
- Hold water molecules together
Dipole: Molecule that have a slightly negative and slightly positive side due to presence of polar covalent bonds
- Polar molecules exhibit permanent role (polarity is permanent).
- Can be identified by identification of polar covalent bonds, which can be identified through comparison of the relative electronegativities of the elements involved in the bond
- Can be identified through analysis of Lewis structures and VSEPR theory.
- Second weakest of the IMFs.
Molecular Solids
Polarity and Polar Molecules
- Polarity of a molecule is determined by combination of the electronegativity of its atoms and its structure.
- Attached by IMFs
- Forces holding molecular solids together are weaker than the forces that held network and ionic solids.
- Low phase change
- Made up of covalent bonds
- Relative softness.
- Relatively low density
- Electrical insulation.
- Because of their IMF's weakness, many molecular solids are liquid or gas at room temperature.
Covalent Bonds and Solids
Covalent bonds form whenever the electronegativity difference between the two atoms involved is not enough to allow one of them to completely take the electron from the other.
- Two types of covalent bonds: Polar & Nonpolar
- Polar covalent bonds: Occur when the electronegativity difference between the two atoms involved is not strong enough for one atom to remove an electron, but strong enough to attract them more to its side.
- Results in slightly positive and slightly negative side to the resultant molecule.
Covalent bonds exist between two nonmetals.
Nomenclature (Covalent Bonds)
Breaking the Octet Rule
Octet rule: All atoms involved in a covalent moelcule must have total of eight electrons.
- Most molecules obey the octet rule; however, some will break the rule like beryllium, born, and nitrogen. They break it because they contain less than eight valence electrons when playing the role of a central atom.
- Any atom in period 3 or beyond is capable of breaking the octet rule because they accept more than eight electrons.
- D-orbital is needed to break the octet rule.
Covalent Nomenclature: Set of rules used to name covalent bonded molecules.
Basic & Acid Nomenclature
- Basic: State the amount of atoms in each element are present with the correct Greek prefix
- Mono = 1, Di = 2, Tri = 3, Tetra = 4, Penta = 5, Hexa = 6, Hepta = 7, Octa = 8, Nona = 9, Deca = 10
- Element with the anion role will have the -ide ending
- Never use mono as first term
- Ending vowel for mono and penta through deca drops when attached to oxygen.
- Acid: First determine if it is a monoatomic or polyatomic acid.
- Starts with H
- Monoatomic Acid: Use the root word of the anion and the prefix hydro- and suffix -ic.
- Polyatomic Acid: Determine whether the polyatomic ion involved ends with -ite or -ate.
- If ends in -ite, change it to -ous.
- If ends with -ate, change it to -ic
- Don't use hypo- for polyatomic acids.
Constructing Lewis Structures
Resonance Structures & Formal Charge
VSEPR Theory
Sometimes, possible to create more than one acceptable Lewis structures for given molecule.
- Formal charge is used in order to determine which of the structures is the correct one.
- Formal charge: Comparison of the electrons present on the valence shell of an atom before bonding and the amount of electrons that still belong to an atom's valence shell after bonding.
- Lowest formal charged structure is correct.
Covalent bonds are represented by Lewis structures.
- Dots are used to represent valence electrons present
- Lines represent covalent bonds between atoms.
- Atoms involved in covalent bonds must have total of eight valence electrons around it when the molecule is complete.
Valence electrons on atoms in a molecule will cause the molecule to arrange itself to minimize electron-electron repulsion.
- Used to predict molecular geometry.
- Most affected by lone pairs
Electron Pair Valence Shell Repulsion:
- Shape it becomes is dependent on number of lone electrons present.
The Quantum Mechanical Model of the Atom
Basics of Atomic Structure
The Periodic Table
The periodic table is arranged in order to demonstrate many patterns present in elements.
- Primarily organized by atomic number.
- Elements are arranged by its period (row) and group (column).
- Groups = "families"
- Elements in the same group show similar levels of reactivity.
- Also organized into metals and nonmetals, solids, liquids, and gasses.
f-orbital
Bohr's model of the atom consists energy levels called "shells".
- The first energy level can hold up to two electrons
- The rest can hold up to eight.
- The energy level the electrons reside in correlate to that element's position on the periodic table.
- Core electrons: Any electrons that are present in the inner shells of an atom
- Valence electrons: Any electrons present on the outermost shell of an atom.
- Electrons live on energy levels. However, they are further subdivided into orbitals and orientations. There are four types of orbitals: s, p, d, and f orbitals.
- Each orientation of a given orbital is capable of holding two electrons with opposing spin; one with positive spin and one with negative spin.
- Exists on energy level after the third.
- Exists in seven different orientations
d-orbital
- Looks like two overlapped p-orbitals
- Exists on every energy level after the second
- Exists in five different orientations
p-orbital
- Shaped like a peanut
- Exists on every energy level after the first
- Exists in three different orientations.
s-orbital
- Shaped like a sphere
- Exists on all energy levels
- Each energy level contains its own s-orbital
- Exists in only one orientation.
Ions
Percent Abundance
Protons and Atomic Number
Atomic Mass
- Atomic number of an element = Amount of protons
- Element of an atom is determined by the number of protons
- CANNOT change the number of protons without changing the element itself.
- Atomic mass is determined by adding protons and neutrons together
- Electrons are not considered because they are so small.
- Allows us to quickly calculate the average mass of an element based on the number of each isotope.
- The greater the percentage, the closer the average value will be to that isotope's mass.
Ions are atoms with charge
- If it is just "atom", then there is no charge.
There are two types of ions:
- Cations = +Charged
- Anions = -Charged
To find the charge...
- Protons - Electrons = Charge
Electronic Configuration
Representations of the quantum mechanical structure of a given atom. There are three types:
Valence Electron Trend
Noble Gas Notation
Valence electron(s) stay the same going down a group and increase by one going across
- Abbreviated form of written electronic configuration.
- Only shows orientation of outermost electrons of a given atom
Atomic Radius
Electronic Configuration Diagrams
The size of the atom
Atomic radius increases going down because it gains a new shell or energy level.
Atomic radius gets smaller left to right because when more protons are added, they attract more electrons.
- Make use of series of boxes and half arrows in order to show how electrons are organized around an atom.
- Shows where atoms are located in terms of energy level, orbital, orientation, and spin.
- Fill from lowest energy to highest energy
Written Electronic Configuration
Ionization Energy
Emission Spectra
- Abbreviated way to show placement of electrons on an atom.
- Shows energy levels and orbitals as well as how many electrons are present on those orbitals, but doesn't directly show orientation and spin.
The amount of energy required to remove an electron from an atom. The closer an atom is to having a full valence shell, the harder it will be to remove an electron from it.
Increases across period because its gets harder to remove an electron as the atom gets closer to filling its valence shell.
Decreases going down a group because electrons get easier to take away the further they are from the nucleus.
Electronegativity
The atom's ability to attract or pull in electrons from other atoms.
Electronegativity increases across period due to atom's increase want of electrons. Increases as atom gets closer to finish its valence.
Electronegativity decrease as electrons get further away. Decreases down a group
An emission spectra for a given element is the different colors that are released when that element is energized.
All light exhibit wave-like properties
- Speed of light = c
- c=3.0x10^8
- Wavelength = denoted by (lambda)
- Frequency = denoted by (nu)
- c = lambda x nu
Quantum Numbers
Ionic Bonds and Ionic Compounds
Atoms will create bonds with other atoms to complete their valence shell.
- Ionic bonds exist when ions are present. Created whenever an anion comes into proximity with a cation.
- Ionic compounds are formed through ionic bonds to completely fill the valence shell of all atoms involved.
There are four quantum numbers: n, l, ml, ms
- n = Energy level
- l = Orbital number
- s=0, p=1, d=2, f=3
- ml = The orientation
- ms = the spin
Ionic Solids
All atoms want to be stable.
- When forming ionic compounds, ions will try to arrange themselves in order to maximize attractive force and minimize repulsive force.
- The chemical formula DOES NOT tell how the atoms are attached; only tells the ratio of the elements' presence.
Characteristics of ionic solids:
- Brittle
- Hardness
- High boiling point
- Electronegativity
Nomenclature
Term that refers to the naming of molecules in chemistry.
- When naming ionic compounds, must provide the entire name of the cation and the root name of the anion with the suffix "ide".
- sodium chloride
- If the element is not in groups IA, IIA, or IIIA, denote its charge with roman numerals.
- CuSO2 = copper (IV) oxide (Cu has a +4 charge, so it is written as copper IV oxide)
History of Atomic Theory
Polyatomic Ions
Polyatomic ions: Ions made of combination of elements
- Oxyanions: Polyatomics made with oxygen. Uses combo of ite, ate, hypo, and per.
- When building compounds with atomic ions, balance the charges.
- Polyatomic ions can exhibit different oxidation states for the elements that make it up
Ex: Iron (III) Sulfate-
Fe=+3 SO4^=-2
3 x 2 = 6
-2 x 3 = -6
0
Fe2(SO4)3
Oxidation States
Oxidation States: Charge of the atom in a molecule or compound.
- Atoms in their natural state have an oxidation state of zero.
- Each atom in a polyatomic ion has its own oxidation state
- The more electronegative anion is the anion, which has a predictable charge
Early 1800s, scientist John Dalton published his concept of the atomic theory. According to Dalton:
- All substances are made of atoms
- Atoms are the smallest particles in existence.
- Atoms of the same element are identical. Atoms of different elements are entirely different
- Molecules of a given chemical compound will always show the same ratio of elements, aka Law of Constant Composition.
- Matter cannot be created nor destroyed.
However, "atoms are the smallest particles in existence" have been proved wrong because the existence of electrons, ions, and isotopes.
J. J Thomson
Ernest Rutherford
- Conducted the Cathode Ray experiment.
- He was able to prove the existence of electrons.
- Proposed the Plum-Pudding model of the atom
- According to the model, the atom is composed of negative charges contained with a positive energy field.
- In effort to prove Thomson's Plum-Pudding model, Rutherford conducted the gold foil experiment. He fired positively charged alpha particles at a thin sheet of gold.
- Small amounts of positively charged particles bounced back. He was able to prove the existence of the nucleus.
- Because the particles deflected were positive, he was also able to prove the existence of protons
Properties of Matter
Physical
Chemical
- Physical properties describe a substance by itself like its shape, size, weight, or color.
- Physical changes are changes in appearance but not the substance or chemical structure
VS
- Chemical properties describe how two things react with each other like flammability, toxicity, or reactivity.
- Chemical changes are changes that alters the chemical structure.
- Some chemical changes are reversible
Separation Techniques
Three ways to to separate solutions:
Chromatography
Distillation
Used to separate two liquids by taking advantages of the substances' different boiling points.
Used to separate a substance into its different parts based on each part's attraction to a medium.
Filtration
Used to separate a solid from liquid through the use of a medium
Mixtures and Solutions
Density
There are two types of mixtures: heterogeneous mixtures and homogeneous mixtures.
- Volume of water= 1 milliliter
- Density is mass over volume
- To find a solid's volume, LxWxH
- The denser an object, the closer the particles are.
Heterogeneous
Homogeneous
Made up of substances where different substances are not evenly distributed.
Examples include:
- Dirt
- worn out brownies
- sushi
- pizza
Made up of substances where different substances are not evenly distributed.
Examples include:
What is Matter?
- Matter is anything that takes up space.
- All matter is made up off small units called atoms
- An element is a type or category of an atom.
- A compound contains more than one element.
- They are formed when two or more different elements are combined.
- There are three types of matter.
Gasses
Liquids
Solids
- Composed of freely moving atoms or molecules
- Exhibits indefinite volume and indefinite shape
- Composed of tightly arranged atoms or molecules
- Exhibits definite volume and definite shape
- Composed of closely arranged atoms or molecules that fluidly move across each other
- Exhibits definite volume, but indefinite shape.