Chapter 5-8 Flashcards
Scientific Law
A major empirical concept that is based on a large body of empirical knowledge
Scientific theory
Well substantiated explanation of some aspect of the natural world that is acquired through the scientific method
Structural Formula
A representation of a molecule in which each covalent bond is shown as a single, double, or triple line between symbols representing bonding atoms
Ex: H-N-H
Ionic bonding
Attraction force holding ions together, simultaneous attraction between positive and negative ions resulting from one or more valence electrons
Ex: Na + Cl > Na+ + Cl- > NaCl
Valence Electrons
An electron in the highest energy level of the atom, an electron available for a covalent bond or electron exchange
Covalent Bond
Simultaneous attraction of 2 nuclei for a shared pair of bonding electrons
Normally forms between 2 nonmetals
Products are molecular substances
Orbital
Specific volume of space in which an electron of certain energy is likely to be found
May contain 2,1 or no electrons
Electrons “spread out” to occupy empty orbitals before forming electron pairs
Valence orbitals
Volumes of space that can be occupied by electrons in an atoms highest energy level
Bonding electrons
A single unpaired electron in a valence orbital that can be shared or exchanged with another atom
Lone Pair
2 valence electrons occupying the same orbital
First Energy Level
One orbital and 2 electrons
Other Energy Levels
4 orbitals and 8 electrons maximum
Octet Rule
Idea that a maximum of 8 electrons (4 pairs) can occupy orbitals in the valence level of an atom
Lewis Symbol
Simple model of the arrangement of valence electrons in atoms or monatomic ions in which the symbol for the elements represents the nucleus plus all the inner electrons and a single or paired dots represent valence electrons
How to Draw a Lewis Symbol
- Write the element Symbol (represent nucleus and bag filled energy levels)
- start by placing valence electrons singly into each four valence orbitals
- If additional locations are required for electrons once each orbital is half filled, start adding 2nd dots
Electronegativity
Describe the relative ability of an atom to attract a pair of bonding electrons in its valence level. Metals have low electronegativity, non-metals have high electronegativity’s
Ex: fluorine has the highest electronegativity being 4.0
Ionic Bonding
Refers generally only to the attraction between any specific cation and any specific anion
- after electron transfers ions arrange themselves in positions where the maximum total attraction between positive and negative occurs> determines the ratio
- arranges in a regular repeating 3D pattern (crystal lattice)
Metallic Bonding
Great number of positive ions surrounded by a sea of mobile electrons, where the valence electrons act like a glue that holds the whole structure together
*do not have to be in any particular arrangement to attract each other> allows for metals to be made into any convenient shape (malleable and ductile)
Molecular Elements
H2(g), N2(g), O2(g), F2(g), Cl2(g), I2(g), Br2(g), P4(s), S8(s)
Structural Formulas
Cl-Cl single
O=O double
Lone pairs are not indicated, each shared pair of bonding electrons is represented by a line, double bond is 2 lines, triple bond is 3 lines
Molecular compounds
Cannot usually be represented by a simplest ratio formula to distinguish between compounds it is necessary to represent them with molecular formulas (accurately represent actual composition of the smallest units of molecular compounds)
Empirical Formula of Water
H2O(l), shows a single molecule contains 2 hydrogen atoms and one oxygen atom held together by covalent bonds
Double Covalent Bond
Explains empirically known molecular formulas such as O2(g), atoms share more than one bonding electron, double bond involves sharing of 2 pairs of electrons between 2 atoms, triple covalent bond involves sharing 2 atoms sharing 3 pairs of electrons
*oxygen, carbon, nitrogen can form more than one kind of covalent bond
Bonding Capacity
Maximum number of single covalent bonds that an atom can form
Ex: carbon has a bonding capacity of 4
Coordinate Covalent Bonds
A covalent bond in which one of the atoms donated both electrons
Ionic Compound Formulas
All have empirical formulas representing crystal lattice with a formula unit, shows only simplest number ratio of cations to anions
- ion charges are omitted and must be memorized
- molar mass: mass of a mole of a formula unit
Empirical Formula
CH2O, shows simplest whole #ratio of atoms in the compound
Molecular Formula
C2H4O2 or CH3OOH, shows actual number of atoms that are covalently bonded to make up each molecule, often has symbols written in a sequence that helps determine which atoms are bonded to which
Stereochemical Formula
3D structural formula, represents molecular shape
Central Atom
Atom to which all other atoms (peripheral atom) are bonded, usually the one with highest bonding capacity
Valence Bond Theory
Successfully explained many of the atomic orientations in molecules and ions
Stereochemistry
The study of 3D spatial configuration of molecular and how this affects their reactions
VSPER theory
Valence-shell-elections-pair-repulsion theory
Based on electrical repulsion of bonded and unbonded electron pairs in a molecule, pairs of electrons in the valence shell of an atom stay as far apart as possible because of the repulsion of their negative charges, electron pairs can be counted by adding bonded atoms and lone pairs
VSPER THEORY STATES
- only valence electrons of the central atom are important for molecular shape
- valence electrons are paired in a molecule or polyatomic ion
- bonded pairs and lone pairs are treated equally
- valence electron pairs repel each other electrostatic-ally
- molecular shape is determined by positions of electrons pairs when maximum distance apart
Linear
AX2
2 bonds and no lone pairs, move to opposite sides.
H–Be–H
Trigonal Planar
AX3, 3 bonds and 3 pairs of electrons around central atom
H
|
B
/ \
H HTetrahedral
AX4, four bonds, four pairs of electrons repelling around central atom
H
|
C
; ¥ \
H H HTrigonal Pyramidal
AX3E, 3 bonding pairs and 1 lone pair
N
; ¥ \
H H H
Angular (Vshaped)
AX2E2, 2 bonding pairs and 2 lone pairs
O
/ \
H H
Linear Pt2 (tetrahedral arrangement)
AXE3, 1 bond pair and 3 lone pairs
H—F
Multiple Bond in VSPER models
Molecule can contain a double covalent bond or triple covalent bond, react rapidly, shorter and stronger
Polar Molecule
One in which the negative electron charge is not distributed symmetrically among atoms making up the molecule.
Partial positive + negative charges on opposite sides
Non-polar Molecules
Molecule with symmetrical electron distribution
Polar Molecule Rules
AB- diatomic different atoms (HCl)
NxAy- nitrogen +other atoms (NH3)
OxAy- oxygen + other atoms (H2O)
CxAyBz- carbon + 2 other atoms (CHCl3)
Non-polar rules
Ax- all elements (Cl2)
CxAy- carbon + 1 other atom (CO2)
Electronegativity rules
Increases with atoms farther to the right or higher up in a column in the periodic table
Non-polar Covalent
When bonded atoms have same electronegativity will attract any shared electrons equally
Polar Covalent Bond
If atoms have different electronegativitys, great the electronegativity difference the more polar the bond.
Can transfer 1 or more electrons to form cations and anions will group to form ionic compound with ionic bonding
Partially negative
The Electrons that spend more time closer to one atomic nucleus than the other spend more time other = partially positive
Bond dipole
The charge separation that occurs when the electronegativity difference of 2 bonded atoms shifts the shared electrons making one end partially positive and one end partially negative
Polar Substances
Polar compounds are soluble in polar solvents=water, non-polar compounds are soluble in nonpolar solvents= grease and oils
Mixing nonpolar and polar substances will form layers with least dense liquid on top
*polar molecules attract more strongly and stay closer together this excludes nonpolar molecules
Intermolecular forces
Forces of attraction and repulsion between molecules
Van der Waals forces
Weak attractive forces that molecules exert on each other (includes London + dipole>dipole forces)
Much weaker than covalent bonds, responsible for what we physically observe about molecular substances, control physical behaviour (boiling point etc)
Dipole-dipole force
Attraction between dipoles, due to simultaneous attraction between any dipole and surround dipoles
Extremely weak, have pronounced effect on ability of solvents to dissolve some solutes and not others
*strength depends on overall polarity of molecule
London Forces (dispersion forces)
Simultaneous attraction between a momentary dipole in surrounding molecules, strength depends on #of electrons and distance between molecules
Predicting Boiling point
As number of electrons increases, boiling point increases
Count electrons as an indicator of relative London force strength
Isoelectronic Molecules
Molecules with the same number of electrons, predicted to have same strengths for London force or intermolecular future(provided shapes are similar)
More polar Molecule=stronger dipole-dipole force =higher boiling point
Greater number of electrons =higher boiling point
Hydrogen bonds
A hydrogen nucleus (a proton) could be shared between pairs of electrons on adjacent molecules results in strong intermolecular force.
- Hydrogen atom must be covalently bonded to another vey electronegative atom
- At least one lone pair of electrons on atoms bonded to hydrogen
* any OH and NH bond
Physical properties of liquids
Surface tension=acts like elastic skin (attraction downwards and sideways)
Meniscus+capillary action= attraction between like molecules (cohesion), and attraction between unlike molecules (adhesion)
Volatility= how rapidly it tends to evaporate
Ionic Crystals
Ionic compounds with low solubility= rocks and minerals, ionic compounds are how reactive metals are found
Formation of ionic compounds
Solid at SATP, held together in rigid structure involves collisions between metal and nonmetal atoms that results in transfer of electrons >cations + anions with filled valence energy levels
Binary ionic compound
Reaction of metal + nonmetal
Crystal Lattice
Arrangement of ions for a given compound, brittle due to ions cannot rearrange without breaking ordered structure of crystal Lattice apart
Simple structure, medium to high melting + boiling
Cubic Fracture
Break along surfaces at 90 degrees to each other
Sodium carbonate
Covalent bonds in each polyatomic ion, holds together as a group, cause group to act like monatomic ion
Complex, varying hardness and melting/boiling points
Non directional
All positive ions attract all nearby negative ions chemical formula, shows only a formula unit expressing the simplest whole number ratio of ions
Ionic crystals
All one structure
Metallic Crystals
Shiny, silvery, flexible solids with good electrical and thermal conductivity.
Soft to hard, melting points are low to high
Continuous very compact crystalline structure, metallic bonding, low electronegativity, empty valence orbitals, electrostatic attractions
Molecular Crystals
Solid Crystals with low melting point, not very hard, non conductive, regular Lattice and packed close, entities are neutral
Covalent Network
(Diamond and quartz), very hard and brittle, high melting points, insoluble, non-conductors of electricity. Do not bend under pressure and seldom break
Covalent Network
Network of covalent bonds, continuous bonding held every entity (atom) in the crystal together, very strong
Carbon
3-D tetrahedral arrangements, layers of sheets, 60-atom spherical molecules, long thin tubes
Semiconductors
Solid state sandwich of crystalline semiconductors used in transitors, Control electrical properties of covalent crystal to produce desired conductive properties
Highest energy levels are full of electrons unable to move, require little energy to jump from to next higher level
