Flashcards in Chemistry 11 Unit 6 Matter, Chemical Trends and Chemical Bonding Deck (31):
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Nuclear Atom
-Properties of subatomic particles (protons, neutrons, and electrons)
Protons and neutrons located in nucleus; electrons located on outer orbiting shells/energy levels
Protons: mass 1 (ratio), + charge
Neutrons: mass 1, neutral charge
Electrons: mass 1/1840 (very small, little mass), - charge
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Atomic number, mass number
Atomic number (Z): number of protons in the nucleus. In a neutral atom, the atomic number is also the number of electrons (PEA)
-protons determine the identity of the element (rise of isotopes that have different number of neutrons but same number of protons)
Mass number (A): sum of the protons and neutrons in the atom
General Notation: opposite of on periodic table
A
X(element)
Z
IONS: charged atoms
+ charge = lost electron (more +)
- charge = gain electron (more -)
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Isotopes
atoms with the same number of protons, but different number of neutrons (therefore same element)
Notation:
A
X(element)
Z
*Z all the same; A different
-isotopes have identical chemical properties (meaning they all react the same)
-Differ in physical properties due to different in mass (i.e. boiling, melting points)
*radioisotopes: radioactive and emit radiation
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Calculating Average Atomic mass from %abundance of isotopes
non-integer atomic masses (i.e. 24.3g/mol) is due to isotopes
average atomic mass (Given in booklet):
sum(mass of isotope x %abundance)/100
*%abundance NOT as decimal
ex: (mass x %) + (mass x %)...
*mass in g/mol or amu (atomic mass unit)
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Iso-Electronic Species
atoms that contain the same number of electrons
I.e. O2- and Mg2+ are isoelectronic (both have 10 electrons)
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Electronic Configurations
describing the electron arrangement/configuration of atoms in terms of main energy levels
Typically can hold max 2, 8, 8
Valence shell: outer (highest) energy level/shell -->important for bonding and chemical reactivity
Ex: Oxygen (Z = 8) electronic configuration 2, 6
*Watch out for charged ions (configuration changes due to gain or loss of electrons)
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Determining relationships between particles
If protons same = same element
If electrons different (protons and neutrons same) = ions (charged)
If neutrons different (protons same) = isotopes
If protons and neutrons different = different elements
If protons and neutrons different (but electrons same) = iso-electronic
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Period Trends (4)
Atomic/Ionic Radius-half the distance between the nuclei of two bonded atoms; the distance from the centre of an atom to its outermost edge
Ionization Energy- energy required to remove an electron
Electron Affinity-energy released
Electronegetivity-covalent bonds
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Period Trends
-Zeff
Across a Period: use Zeff
-increase in protons and electrons in same energy level, increase attraction
Down a Group: use energy levels
-greater distance, increase in core electrons/shielding electrons, less attraction
Zeff: Effective nuclear charge
-attractive force between the protons in the nucleus and the valence electrons in an atom
Zeff = Z - S (# of core electrons/shielding effect)
-equal to group # ex: Zeff of fluorine is 7
*Shielding electrons cancel out attractive force
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Energy levels increase going down a group
Explanation
Down the group, electrons are being added to increasingly higher energy levels at a greater average distance from the nucleus
Result of More energy levels
-larger atom
-decrease in attraction
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Zeff increases across a period
Explanation
across a period, there is an increase in the number of protons and valence electrons in the same energy levels. This results in a stronger attractive force between the protons and the valence electrons
-number of energy levels do no change across a period, so there is no shielding effect to cancel out the attractive force
Zeff remains constant down a group because the increase in the nuclear charge is cancelled out by the increase in shielding electrons (therefore zeff remains the same)
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Atomic and Ionic Radius
Decreases across a period: Zeff increases
-increasing number of protons and an increasing number of valence electrons in the SAME energy level
-attraction pulls energy levels closer to the nucleus, thus shrinking the atom
Increases down a group: energy levels
-valence electrons are being added to increasingly higher energy levels at a greater average distance from the nucleus
-also increased electron-electron repulsion among inner electrons (due to more energy levels holding more electrons)
-result in a reduce of attraction between the positive nucleus and the valence electrons, causing the atomic radius to increase
Ionic Radi refers to radius of ions (charged atoms)
-atoms lose electrons (positively charged cations) = smaller than parent atom -->minimizing electron-electron repulsions (making atom smaller)
-atoms gain electrons (negatively charged anions) = larger than parent atom -->increased electron-electron repulsion which tend to push electrons away from each other (making atom larger)
*Overall trend is still the same (going down larger, across smaller)
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Ionization Energy
-in metals (ionic bonding) because metals tend to lose electrons
Refers to energy required (endothermic) to remove an electron from an atom in the gas phase.
-also indicates how strongly an atom's nucleus holds on to its valence electrons
Greater ionization energy, more difficult to remove
-non-metals have high ionization energy (want to receive electron, not lose)
-metals have low ionization energy (want to lose electrons due to few number of valence electrons)
Increases across a period: Zeff increases
-increase number of protons and electrons in the same energy level
-increased attraction between electrons and nucleus that makes it more difficult to remove an electron (need more energy)
Decreases down a group: energy levels
-increasing distance between valence electrons and nucleus
-weaker attraction, easier to remove electron
*typically asking for first ionization energy (energy required to remove the first electron)
-more energy is needed to remove each successive electron due to electrons being removed from a stable outer shell
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Electron Affinity
the release of energy (exothermic) when electrons attach to the atom
-also attraction an atom has for electrons other than its own
-specific for non-metals gaining electrons in ionic bonding
-non-metals have high electron affinity (want to gain electrons)
**DO NOT look at sign - or + on chart (compare values)
Increases across a period: Zeff increases
-as protons and electrons increase, more strongly attract electrons from other atoms (non-metals want to complete their octet)
Decreases down a group: Energy levels
-electrons are being added to increasingly higher energy levels at greater average distance from the nucleus
-large distance results in less attractive force to add/gain electrons
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Electronegativity
-covalent bonding
-attraction an atom has for electrons in a covalent bond
-ability of an atom in a molecule to attract a shared electron to itself
Increases across a period: Zeff increases
-increasing nuclear charge generates greater attractive forces on shared bonding electrons
Decreases down group: energy levels
-increasing distance between nucleus and shared electrons
-since the shared electrons are in increasingly higher energy levels at greater average distance from the nucleus, there is less attraction for the shared bonding electrons
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Types of Bonding
Ionic Bonds: metal and nonmental (positive and negative charges)
Covalent Bonds: non-metal and non-mental (sharing of electrons)
Metallic Bonds: bonding electrons are free to move throughout the 3 dimensional structure (mobile sea of electrons) -->conductive (ex: alloys)
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Valence Electrons
electrons residing in the outer shell, electrons that are involved in chemical interactions and bonding
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electron-dot (lewis) symbols for elements
-convenient representation of valence electrons
-allows you to keep track of valence electrons during bond formation
-consists of the chemical symbol for the element plus a dot of each valence electron
-each side of the element can accommodate up to 2 electrons
-the number of valence electrons is the same as the column number of the element in the periodic table
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Octet Rule
atoms tend to lose, gain or share electrons until they are surrounded by 8 valence electrons
-want to end up with 8 valence electrons (stable)
*Exceptions (i.e. He and H = duplet rule)
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Ionic Bonding
-ions are arranged in a crystal lattice structure
-opposite charges
Represented using Lewis-dot symbols
Element 1 (with dots) + element 2 (with dots or other shape) --------> [element 1]+charge[element 2]-charge
**Need to show electron transfer (using arrow)
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Covalent Bonding
-sharing electrons to achieve the configuration as noble gases
Lewis Structures: 2d model that represents covalent bonds as lines and the unshared electrons as dots
-does not show inner electrons
One pair of electrons shared: single bond
Two pairs of electrons shared: double bond
Three pairs of electrons shared: triple bond
*the distance between the bonded atoms decreases as the number of shared electron pairs increases
i.e. triple bonds are much shorter and therefore have a stronger bond than single bonds
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Bond Polarity and Electronegativity in Covalent Bonds
-how to determine type of bond
Extreme Examples
-equal sharing of electrons (same atom)
-out electrons are stripped from on atom
Bond Polarity: describes the sharing of electrons between atoms
-Nonpolar covalent bond: electrons are shared EQUALLY between two atoms i.e F2
-Polar covalent bond: one atom has a greater attraction for the electrons than the other atom i.e. HF
Electronegativity: quantity used to determine whether a given bond will be non polar covalent, polar covalent, or ionic
-defined as the ability of an atom in a particular molecule to attract electrons to itself (the greater the value, the greater the attractiveness for electrons)
-relates to ionization energy and electron affinity
Highly electronegative atoms: attract electrons but resist having its own electrons taken away
Most electronegative element = Fluorine (4.0)
Least Electronegative element = Cesium
See electronegativity different to determine bond (between two atoms/elements)-->see rules in book
-when the electron is completely stripped from an atom, it is an ionic bond (no sharing)
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Polar bonds (represented as)
arrow facing towards more electronegative atom
+--------->
delta + delta -
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Drawing Lewis Structures for Covalent Compounds
1) Calculate total number of bonds
Total - Valence (+/- charge)/2
2) draw dot/line structure (no shape yet)
3) determine # of bonding pairs and non-bonding pairs
4) look at chart to determine shape
-EDG, MG
5) draw finalized structure (ensure each atom has a complete octet)
-INCLUDE BRACKETS for ions
Hints:
-central atom is usually listed first
-H is always terminal, halogens are almost always terminal (single bonds form)
-oxygen often forms double bonds
-carbon forms single, double or triple bonds
-nitrogen forms triple bonds
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Molecular Geometrics
-Lewis structure provides no information about the structure/shape of the molecule
-structure of a molecule is defined mainly by the BOND ANGLES
tri = 120degrees
tetra = 109.5 degrees
if have double bonds or lone pairs, then will be less than (
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VSEPR
Valence Shell Electron Pair Repulsion model
-bonding pairs of shared electrons tend to repel other bonding pairs of electrons in the valence orbital (wants to be far apart)
the best spatial arrangement of the bonding pairs of electrons in the valence orbitals is one in which the REPULSIONS are MINIMIZED
-therefore, we can predict molecular geometrics
Steps in determining VSEPR model:
1. draw Lewis structure
2. count total number of electron pairs around CENTRAL atom (# of bonding and non-bonding pairs). Use data booklet to see best arrangement to minimize the electron shell repulsion
3. describe electron-domain geometry and molecular geometry
*Double or triple bond is counted as ONE bond when predicting geometry
**Draw balloons around lone pairs to signify compressing other atoms around the lone pairs
ex: sometimes think it is linear, but actually tetrahedral (lone pairs on central atom)
The bond angles decrease as the number of non-bonding electron pairs increases
-due to higher electron density
Electrons in multiple bonds (i.e. non-bonding) exert a greater repulsive force on adjacent electron pairs than do single bonds.
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Polarity of molecules
overal "charge distribution" of a molecule is determined by shape and polarity of bonds
Polar Molecule:
-molecule has a dipole (- and + charge on each end)
-not symmetrical (due to similar distribution of charges & electronegativity)
-non-bonding/lone pairs
-uneven distribution of charges
ex: HF, H20
Nonpolar molecule:
-does not have a net dipole
-symmetrical
-even distribution of charges
Ex: Cl2, linear and tetrahedral
if one atom has more force than the other (different atoms), then polar molecule
*NOT all molecules with polar bonds exhibit dipoles (not all polar bonds are polar molecules)
-symmetry of bonds cancel out asymmetry of charges
80% of molecules are polar
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Polarity of Polyatomic Molecules
each polar bond in a polyatomic molecule will have an associated dipole
the overall dipole of the molecule will be the sum of the individual dipoles
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Dipole moment
separation of the charge around the molecule into a more positive and negative area
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Bond Angles
linear = 180
tri = 120 (
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