Unit 2: Structure and Properties of Matter Flashcards
(87 cards)
1
Q
Who discovered the atom?
A
The Greeks
2
Q
Who invented the Modern Scientific Method
A
Robert Boyle
3
Q
Who invented the Atomic Theory of Matter?
A
John Dalton
4
Q
What are the 5 parts of the Atomic Theory of Matter?
A
1) All matter is made up of tiny particles
2) All atoms of a given element are identical
3) Atoms of different elements have the same properties
4) Atoms of different elements combine in constant ratios to form compounds
5) Chemical reactions involve the rearrangement of atoms. Atoms can’t be created or destroyed
5
Q
What did JJ Thomson discover?
A
Atoms of many elements can be made to emit tiny negative particles
6
Q
What is the Plum-Pudding Model (AKA Rasin Bun)
A
Negative electrons that are embedded into a positively chargers and spherical cloud
7
Q
Who performed the gold foil experiment?
A
Ernest Rutherford
8
Q
Who discovered the Neutron?
A
James Chadwick
9
Q
What did Mac Planck start?
A
The quantum revolution but studying light emitted on hot objects
10
Q
Quanta
A
A series of short burst of energy rather than a consistant steam
11
Q
Quantum
A
One burst or packet of energy
12
Q
Who discovered the photoelectric effect?
A
Heinrich Hertz
13
Q
What is the formula for the photoelectric effect?
A
E=hf
- E: Energy in joules
- h: Planck’s constant (6.6 X 10^34 J/Hz)
- f: Frequency in hertz (Hz)
14
Q
What are the 2 atomic states of hydrogen?
A
1) Excited State; Atom with excess energy
2) Ground State; Atom in lowest possible state
15
Q
When an atom absorbs energy from an outside source what atomic state does it enter?
A
Excited
16
Q
Whats the basis of boar’s theory of the hydrogen atom?
A
He used his observations of the line spectra to explain why electrons dont collapse into the atoms nucleus
17
Q
Energy Level Diagram
A
- Energy in the photon that corresponds to the energy used by the atom to reach the excited state
- Each element produces it own unique colour spectrum which can be used to identify elements in compounds and mixtures
18
Q
Quantized Energy Levels
A
Since only certain energy changes occur the H atom must contain discrete energy levels
19
Q
Whats the 3 main points of Bohr’s model of the atom?
A
1) Quantized energy levels
2) Electron moves in a circular orbit
3) Electron jumps between orbits by absorbing or emitting photon of a particular wave length
20
Q
Principal Quantum Number (n)
A
- Designated the main energy level of an electron
- Can be equal to 1,2,3,4,5,6,7
21
Q
Secondary Quantum Number (l)
A
- Describes the shape of the electron orbit
- Equal to 0,1,2,3
22
Q
Magnetic Quantum(ml)
A
- Describes the orientation of the electron orbit in space
- For each value of l ml can vary from -e to +l
- The number of values for ml is the number of independent orientations of orbits that are possible
23
Q
Spin Quantum Number ms
A
- Describes whether an electron spins clockwise or counterclockwise
- ms can only be +1/2 or -1/2
24
Q
What are the 4 pieces to finding a quantum number set?
A
1) Principal Quantum Number (n)
2) Secondary Quantum Number (l)
3) Magnetic Quantum(ml)
4) Spin Quantum Number ms
25
Energy Level Diagram
Shows the relative energies of electrons in various orbitals
26
What are the 6 rules for filling energy level diagrams?
1) Each circle is an orbital that can hold up to 2 electrons
2) Electrons are represented by arrows (# of arrows will equal atomic number of atom)
3) Arrows point up or down depending on their spin
4) s=1 circle (2 electrons)
5) p=3 circles (6 electrons)
6) d=5 circles (10 electrons)
27
Pauli Exclusion Principal
No 2 electrons in an atom can have the same 4 quantum numbers, meaning that the arrows in the orbitals must point in different directions
28
Aufbau Principle
Each electron is added to the lowest energy orbital available
29
Hund's Rule
One electron occupies each of the orbitals at the same energy before a second electron can occupy the same orbital
30
Electron Configuration
- Provides the same information as the energy-level diagrams but in a more concise format
- Lists the number and kinds of electrons in order of increasing energy, written in a single line
31
What are the pros and cons of using Electron Configurations?
Pros- Efficient
| Cons- Some energy may be lost
32
Shorthand Electron Configurations
Same as regular configurations but you start with the noble gas proceeding the element of interest
Example: Cl
[Ne] 3s^2 3p^5
33
What are the 4 rules for drawing lewis structures?
1) Use the last digit of the group on the periodic table to determine the number of valence electrons
2) Place one electron on each of the 4 sides of the nucleus before pairing
3) If there are more than 4 valence electrons pair them tas required
4) Use unpaired electrons to bond additional atoms with unpaired electrons to the central atom
34
What are the 6 steps to the procedure for drawing Lewis Structures? (Look in loose notes for examples)
1) Arrange atoms symmetrically around the central atom; usually listed first in the formula
2) Count the number of valence electrons of all atoms. For polyatomic ions, add electrons corresponding to the negative charge and subtract electrons corresponding to he positive charge on the ion
3) Place a bonding pair of electrons between the central atom and each surrounding atom
4) Complete the octets of the surrounding atoms using lone pairs of electrons (Hydrogen is an acception) Any remaining atoms go on the central atom
5) If the central atom doesn't have an octet move lone pairs from the surrounding atoms to form double or tripe bonds until the central atom has a full octet
6) Draw the lewis structure and enclose polyatomic ions within square brackets showing the ion charge
35
Octet Rule
When the outer most shell is full with electrons
36
Valence Bond Theory
A covalent bond is formed when two orbitals overlap (Share the same space) to produce a new orbital containing 2 electrons of opposite spin
37
What are the 4 summary points of Valence Bond Theory?
1) A half-filled orbital in one atom can overlap with another half-filled orbital of a second atom to form a new bonding orbital
2) The new bonding orbital from the overlap of atomic orbitals contains a pair of electrons with opposite spin
3) The total number of electrons in the bonding orbital must be 2
4) When atoms bond they rearrange themselves in space to achieve the maximum overlap of their half-filled orbitals. Maximum overlap produces a bonding orbital of lowest energy
38
What are the 2 problems with the Lewis Bonding Theory?
1) The 4 equal bonds represented by the 4 pairs of electrons in a carbon compound ( Example: Methane, CH4)
2) The existence of double and triple bonds
39
Hybridization
A theoretical process involving the combination to create a new set of orbitals that take part in covalent bonding
40
Hybrid Orbital
An atomic orbital obtained by combing at lease two different orbitals
41
Sigma Bond; σ
A bond created by end-to-end overlap of atomic orbitals
42
Pi Bond; π
A bond created by the side-by-side (or parallel) overlap of atomic orbitals
43
Valence Shell Electron Pair Repulsion Theory (VSPR)
- Based on the electrical repulsion of bonded and bonded electron pairs in a molecule or a polyatomic ion
- The number of electron pairs can be counted by adding the number of bonded atoms plus the number of lone pairs of electrons
44
What are the 5 Points of summary for the VSPR Theory?
1) Only the valence shell electrons of the central atom are important for molecular shape
2) Valence shell electrons are paired
3) Bonded pairs of electrons and lone pairs of electrons are treated approximately equally
4) Valence shell electron pairs repel each other electro statically
5) The molecular shape is determined by the position of the electron pairs when they are a maximum distance apart ( With the lowest repulsion possible)
45
What are the 2 steps to determining the shape of a molecule using the VSPR Theory?
1) Draw the lewis structure for the molecule, including the electron pairs around the central atom
2) Count the total number of bonding pairs (Bonding atoms) and lone pairs of electrons around the central atom
46
Polar Molecules
A bond that results from a difference in electronegativity between the bonding atoms; one end is - and one end is +
47
What is the relationship between electronegativity and polarity of a chemical bond?
The greater the difference in electronegativity the more polar the bond
48
A very polar bond is an...
A non-polar bond is a...
A somewhat polar bond is a...
Ionic bond
Covalent bond
Polar covalent bond
49
Non-polar bond
Results from a zero difference in electronegativity between the bonded atoms; a covalent bond with equal sharing of bonding electrons
50
Ionic Bond
A bond between a metal and a non-metal
51
Covalent/Non-polar
A bond between 2 non-metals
52
Polar Covalent Bond (Polar Bond)
A bond where electrons are shared somewhat equally
| Example: HCl
53
What is the general rule used to determine if a bond is ionic?
When the difference in electronegativity exceeds 1.7 the percent ionic character exceeds 50%
54
Bond Dipole
The electronegativity difference of 2 bonded atoms represented by an arrow pointing from the lower to the higher electronegativity
55
Non-polar Molecule
A molecule that has either non-polar bonds or polar bonds whose bond dipoles cancel to be zero
56
Polar Molecule
A molecule that has polar bonds with dipoles that dont cancel to zero
57
What is true about symmetrical molecules?
The sum of the bond dipoles is zero and the molecule is non-polar
58
What are the 4 points of summary for Theoretical Prediction of Polarity?
1) Draw a lewis structure for the molecule
2) Use the number of electron pairs and VSEPR rules to determine the shape around each central atom
3) Use electronegativities to determine the polarity of each bond
4) Add the bond dipole vectors to determine if the final result is zero (non-polar or nonzero (polar)
59
VanderWaals Forces
Molecules exert forces on each other
60
What are the 3 types of VanderWaals Forces (Intermolecular Forces)
1) Dipole-Dipole
2) London Forces
3) Hydrogen Bonding
61
Which is weaker...Intermolecular forces or covalent bonds?
Covalent Bonds
62
Dipole-Dipole Forces
- The simultaneous attraction of one dipole by its surrounding dipoles
- The strength of the dipole-dipole force is dependant on the polarity of the molecule
63
London Forces
- Due to the simultaneous attraction of the electrons of one molecule by the positive nuclei in the surrounding molecules
- The strength of the London force is directly related to the number of electrons in the molecule
64
London force theory states what?
As the number of electrons increases, so does its boiling points
65
Isoelectric
Having the same number of electrons per atom, ion or molecule
66
What are the 5 points on predicting with dipole-dipole and London forces?
1) The more polar the molecule the stronger the dipole-dipole and therefore the higher the boiling point
2) The greater the number of electrons per molecule the stronger the London Force and therefore the higher the boiling point
3) If molecules are isoelectric they have the same strength London force
4) If the London forces are the same; you can predict the boiling points of 2 substances by comparing their dipole-dipole force
5) If the dipole-dipole forces are the same; you can predict the boiling point of 2 substances by comparing their London force
67
Hydrogen Bonding
- The attraction of atoms bonded to N, O, or F atoms to a lone pair of electrons of N, O or F atoms in adjacent molecules
- Strongest
68
What 2 things do Intermolecular Forces affect?
1) Surface Tension
| 2) Capillary Action
69
What is trapped in the ice found in the ice found in the arctic?
Methane
70
What are the 4 classes of substances and their combined elements
1) Ionic; Metal and non-metal
2) Metallic; Metal
3) Molecular; Non-metals
4) Covalent Network; Metalloids/carbon
71
Ionic Crystals
A 3-D arrangement of + and - ions held together by strong, directional ionic bonds
72
What are the 3 properties of Ionic compounds?
1) Hard but brittle solids
2) Conduct electricity when dissolved in water
3) High melting points
73
What is the strongest intermolecular force?
Ionic bonding
74
Metallic Crystals
A 3-D arrangement of metal cations held together by strong non-directional bonds created by a "sea" of mobile electrons
75
Molecular Crystals
A 3-D arrangement of neutral molecules held together by relatively weak intermolecular forces
76
What are the 3 properties of molecular crystals?
1) Relatively low melting points
2) Soft
3) Non conclusions of electricity (Neutral molecules)
77
Covalent Network Crystals
A 3-D arrangement of atoms held together by strong , directional covalent bonds
78
What are the 4 properties of covalent network crystals?
1) Very hard and brittle
2) High melting points
3) Insoluble
4) Non-conducting
79
What makes the Carbon-carbon bond structure so strong?
An interlocking structure
80
What are the 4 carbon structures?
1) Diamond
2) Graphite
3) Buckyball
4) Carbon Nanotubes
81
Review l and the periodic table
.
82
Review the energy level diagram
.
83
Review Electron Configurations
.
84
Review Valence Bond Theory
.
85
Review VSEPR Theory
.
86
Review Polar Molecules
.
87
Review Crystalline Solids
.
