The Structure of Matter - St. Johns County School District
The Structure of Matter How Matter Combines and What It May Make Compounds And Mixtures What are compounds? Two or more elements bonded together Chemical bonds distinguish compounds from mixtures The force that holds two atoms together The ability to write a formula, such as H2O, indicates a compound A compound always has the same chemical formula Types of Bonds Ionic: Opposing ions attract one another Covalent: Atoms share e-
Metallic bonds: Metal atoms bonding with other metal atoms Hydrogen: Bonding via hydrogen or polar molecules Bonds can bend and stretch without breaking (bond lengths are averages) Models of Compounds 1. Ball and Stick good for showing angles Caffeine 2. Space-filling
good for relative atomic sizes 3. Structural formula most commonly used Various Hydrocarbons How Does Molecular Structure Affect Properties? Recall that ionization energy refers to how easily an atom loses an electron Electron affinity refers to how much attraction an atom has for electrons
Nobel gases have a full valence shell Thus, they have high ionization energy and low electron affinity THEY ARE INERT!!! Means they dont react or dont want to change Elements tend to react to acquire the stable electron structure of a noble gas Formation of Ions Cation By losing an electron, the atom has gained a positive charge (+1) and now has a different electron configuration Anion Elements on the right side of table have high electron affinity and will therefore accept electrons forming a negative charge (-1) Ionic Bonds
A type of chemical bond that holds oppositely charged particles together in an ionic compound Oxides: Metal and Oxygen (O) Salts: Metal and non-metals Electrons are transferred, not shared Ionic compounds are in the form of networks, not molecules Ionic Bonds During the formation of ionic bonds, the positive and negative ions are packed into a regular repeating pattern that balances the forces This strong + to attraction in ionic bonds create a crystal lattice
This causes all the cations to be completely surrounded by anions and vice versa Properties of Ionic Compounds Crystalline solids STRONG BONDS Hard and brittle High melting points High boiling points High heats of vaporization
High heats of fusion Good conductors of electricity when molten - dissolved Poor conductors of heat and electricity when solid Many are soluble in water Energy of the Ionic Bond Creating ionic bonds is always exothermic exo means out or exit, thermic means thermal or heat
Heat or energy is released If equal amount of energy were inserted into the bond, then the ionic bond would break The more negative the lattice energy, the stronger the force of attraction Larger absolute charges and smaller ions have a larger negative lattice energy Covalent Bonds A chemical bond that results from sharing of valence electrons The shared electrons are considered to be part of the complete energy level of both atoms This is also known as a molecule These tend to form between nonmetals Covalent Properties Covalent bond involves attractive and repulsive forces
With molecules, it is endothermic to break bond and exothermic when building it, same as ionic Covalent bonds can form two different types: Polar and non-polar Covalent Bonds Example of these can be found in the naturally occurring diatomic molecules Remember Mr. BrINClHOF Bromide (Br), Iodine (I), Nitrogen (N), Chlorine (Cl), Hydrogen (H), Oxygen (O) and Fluorine (F) Single Covalent Bonds This occurs when a single pair of electrons is shared The shared electron pair, or bonding pair, is represented by either a pair of dots or a line Lewis Dot Structures Multiple Covalent Bonds Atoms can attain a noble-gas configuration by sharing more than one pair of electrons between two atoms
Mostly formed by C, N, O and S This can be seen with double and triple covalent bonds Chemical Bonds Non-Polar Covalent bonds are between two identical nonmetal atoms Polar Covalent bonds are between two different nonmetal atoms Ionic bonds primarily form between nonmetals and reactive metals Polar vs. Non-polar Recall: Electronegativity is the attraction of electrons via atoms bonded together in molecules If the difference in electronegativity between the bonded atoms is great
enough, then the electrons will orbit more frequently around the atom with greater electronegativity Equal sharing of electrons are non-polar molecules Unequal sharing creates polar molecules Polar Molecules Polar molecules will have the uneven distribution of electrons because of the nucleuss pull This creates POLES to the molecule similar to the Earths North and South Poles e.g.: HF Hydrogen (2.2) vs Fluorine (4.0) Electrons will be found closer to fluorine than hydrogen, creating a negative pole are the fluorine nucleus Thus, a positive pole near hydrogen This created a DIPOLE molecule (two poles) + -
Covalent vs. Ionic Different Different Share electrons Transfer electrons Alike (ions formed) (polar vs. nonpolar) Topic Between Two Nonmetals Weak Bonds
(low melting point) Chemical Bonds Covalent +/- Topic Ionic Electrons are involved Between Metal and Nonmetal Strong Bonds (high melting point) Metallic Bonds
Electrons move freely between metal atoms An attraction between nuclei and neighboring electrons packs the metals in tightly, and allows electrons to move freely between This is why metals conduct electricity Metallic Bonding Cations + e1+ e 1- +
+ e electron sea 1- e1- + e1- + e1- e1+ e1- Metallic bonding is the attraction between positive ions and surrounding freely mobile electrons Most metals contribute more than one mobile electron per atom Shattering an Ionic Crystal; Bending a
Metal broken crystal An ionic crystal Force + + + + - + + + + - + + + + -
+ + + + + + + + + + + + + + + + No electrostatic forces of repulsion metal is deformed (malleable) Chemical Bonds (The Glue) Ionic Bonds What holds ionic bonds together? The positive cation(s) and the negative anion(s) attraction via the electromagnetic force
IMPORTANT: There are four universal forces Strong Force: holds nucleus together (quarks or nucleon glue) Electromagnetic Force: attracts opposite charges and poles Weak Force: allows transmutation of nucleons Gravity: mass attracts mass Chemical Bonds (The Glue) Covalent Bonds Van der Waals Forces What holds a polar covalent molecule together? The partial positive and partial negative charges via the electromagnetic force (again) Also known as dipole forces Then what holds a non-polar covalent molecule together? London Dispersion Forces Electrons are always moving! Constantly moving within their own orbits and around the nuclei of the atoms in the bonds At any point, there will be a congregation of electrons in one area more than another This creates small TEMPORARY dipole forces that last fractions of fractions of a second
Hydrogen Bonding This type occurs when hydrogen has a partial positive charge from the unequal distribution of electrons Electronegativity The partial positive hydrogen will bind with a partial negative atom on another molecule This connects water to water for example Dipole-dipole (two polar molecules binding) Polyatomic Ions Within Compounds Groups of covalently bonded atoms that have either lost or gained electrons Behave as ions (cations or anions) Parentheses group the atoms of a polyatomic ion and show that they act as any elemental ion
Cations are placed in front and anions behind The charge of the polyatomic ion is for the entire group, not just part of it Names may be related to oxygen content Examples of common polyatomic ions: 2- Charge 1- charge Formula Name C2H3O21- acetate
ClO31- chlorate ClO21- chlorite Formula Name CN1- cyanide CO32- carbonate HSO31- hydrogen sulfite CrO42-
Name NO31- nitrate AsO43- arsenate NO21- nitrite C6H5O73- citrate PO43- phosphate Formula Name
PO phosphate NH41+ ammonium 33 1+ charge Compounds with Polyatomic Ions Ba(CO3)2 NH4Cl Ba(OH)2
K(ClO3) Fe(OH)3 Fe(OH)2 Al(OH)3 Al2(SO4)3 Notice the position of the subscripts If within the ( ) then they are a part of the polyatomic ion, if not, then it is stating the number of the polyatomic ions
Al2(SO4)3 3 polyatomic ions Total of 2 Als, 3 Ss, and 12 Os Compounds with Polyatomic Ions Ba (OH)2 The metal (Ba) is held to the polyatomic ion by ionic bonding These nonmetals are held to each other covalently Naming Ionic Compounds: Charge is determined mathematically Overall charge must be zero unless otherwise stated Fe2O3 has iron and oxygen Cation is first and anion is second If there isnt a polyatomic ion, then it ends in ide All cations need a roman numeral unless its charge
is constant i.e. transition metals can have changing charges Each O is -2 (periodic table) Os equal a total of -6, so Fes must equal +6 There are 2 Fes, so each must be +3 Fe2O3 is Iron III Oxide Examples 1. Fe2O Answer: Iron I Oxide First, -14 this side 2. W2O 7 --- W2 O7 Second, +14 this side Third, 2 Ws must total +14, so each one is +7
Answer: Tungsten VII Oxide 3. Hg(CN)2 --- Hg (CN)2 Second, +2 this side First, -2 this side from the polyatomic ion chart Third, 1 Hg must be +2 by itself Answer: Mercury II Cyanide Naming Covalent Compounds Numerical prefixes are used to name covalent compounds of two or more elements Mono- , Di- , Tri- , Tetra- , Penta- , Hexa- , Hepta- , Octa- ,
Nona- , Deca- Elements are listed left to right on the periodic table and ends in ide Silicon dioxide --- SiO2 If the first element is singular, mono may be left off Dihydrogen Oxide --- H2O (water) Nitrogen Trichloride --- NCl3 Naming Basic Acids Look for two characteristics: 1. The compound/molecule will begin with hydrogen, H This will identify it as an acid
2. Look for what is immediately after the hydrogen, If it is a single element, then it is a binary acid If it is a polyatomic ion, then it is an oxyacid Binary Acids: HCl, HBr, HI, etc. Called hydro- name of the anion with an ic ending, followed by acid Hydrochloric acid, hydrobromic acid, hydroiodic acid, etc. Naming Basic Acids An oxyacid is when there is a polyatomic ion after the hydrogen, H2SO4 If the polyatomic ion ends in ite, then the acid will end in ous If the polyatomic ion ends in ate, then the acid will end in a ic ending
HNO3 Nitric Acid HNO2 Nitrous acid VSEPR You should now be familiar with bond types, naming and drawing them in 2D (Lewis Dot Structures) Now we will begin to picture them in 3D 3-D models are based on VSEPR: Valence Shell Electron Pair Repulsion Theory VSEPR The idea behind VSEPR is that covalent bonds and lone pair of electrons spread to the furthest distance they can from one another Covalent bonds consist of electrons and they have the same charge Thus, VSEPR explains why molecules have their shapes
VSEPR Flow Chart VSEPR Flow Chart 1. Draw the Lewis Structure for the compound in question 2. Count the number of things 3. Follow the flow-chart Molecule Shapes Tetrahedral Tetrahedral molecules look like pyramids with four faces and each point corresponds to an atom that's attached to the central atom Bond angles are 109.5 degrees Trigonal pyramidal It's like a tetrahedral molecule, except flatter Bond angles are 107.5 degrees (it's less than tetrahedral molecules because the lone pair shoves the other atoms closer to each other) Trigonal planar It looks like the hood ornament of a Mercedes automobile The bond angles are 120 degrees Molecule Shapes
Bent They look bent Bond angles can be either 118 degrees for molecules with one lone pair or 104.5 degrees for molecules with two lone pairs Linear The atoms in the molecule are in a straight line This can be either because there are only two atoms in the molecule (in which case there is no bond angle, as there need to be three atoms to get a bond angle) or because the three atoms are lined up in a straight line (corresponding to a 180 degree bond angle)
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