Gen Chem Flashcards
(122 cards)
1
Q
Isotope A (95%) mass number is 20, B (0.3%, mass #= 21), C (9.3%, 22 mm). What is the element and about what would be the mass number
A
NEON, because we can see that the majority of A is 20 and therefore mass number will be around 20, less than 22 and less than 21)
2
Q
Beta (-) & Beta (+)
A
Beta (-) : converts a neutron into a proton & emits and electron Beta (+) and electron capture: convert a proton into a neutron
3
Q
Half life:
A
is the time required for an amount of a given isotope to decrease by half –>with each subsequent half-life interval that passes, the amount that remains decreases by half again, becoming exponentially smaller and eventually approaching zero. –>the molar concentration of a solution at a given time is determined by dividing the number of moles analyte by the sample volume (in liters)
4
Q
Density:
A
mass/volume —>materials tend to expand when heated & contract when cooled, the volume that a substance occupies can undergo minute changes with temp –>as a result, the exact density of a substance varies slightly with temp —>atoms packed
5
Q
Bohr Model assumptions
A
- electrons move around the nucleus in fixed circular orbits at particular intervals. 2. Electrons in orbits farther from the nucleus have higher energy 3. Energy equal to the difference between 2 orbits is absorbed by an electron moving to a higher orbit & is emitted by an electron moving to a lower orbit 4. energy that is absorbed or emitted by an electron equals the energy difference between two orbits
6
Q
First Ionization energy
A
energy required to remove the first, most loosely bound valence electron from a neutral atom –>the 1st ionization energy tends to increase with increasing atomic number moving across a period & decrease moving down a group on the periodic table
7
Q
Effective Nuclear Charge
A
Zeff= Z - S Z is nuclear charge S is shielding constant (number of core electrons) Results: an effective nuclear charge that is less than Z & increases as the atomic number increases
8
Q
Elements within the same group:
A
have the same number of valence electrons (similar chemical properties) but can have very different physical properties
9
Q
When getting the electron configuration for an ionzed atom (Ca+2), which electons are lost first?
A
lose electrons from higher energy first (4s)
10
Q
Paramagnetic & Diamagnetic
A
Paramagnetic: configuration WITH unpaired electrons
Diamagnetic: configuration without unpaired electrons
11
Q
_____________ is the extent to which an electron cloud of an atom can be distorted by an external charge or by an applied electric field to produce a dipole
A
Polarizability
12
Q
____________is the tendency of an atom to attract electrons within a bond
A
Electronegativity
13
Q
__________ assesses the tendecny of an atom to accept an additional electron by measuring the energy change when an electron is added to an atom
A
Electron affinity
14
Q
___________ measured the ENERGY required to REMOVE AN ELECTRON from an atom
A
Ionization energy (opposite of electron affinity) –>high IE=least reactive —> reactivity of atoms forming ionic compounds INCREASES as the IE DECREASES (moving down alkaline-earth metal column, IE decreases. This makes removing an electron more favorable and increases reactivity) -valence are removed first, then core electrons -removing a core electron takes MORE ENERGY than does removing a valence electron
15
Q
Second Ionization Energy
A
***ionizations involving core electrons are higher energy than those involving valence electrons **** Example: Na & Mg -Group 1 (Na) 2nd IE is much higher than Mg (group 2) because Na has only once valence electron and removing a second electron from Na requires the loss of a core electron
16
Q
Object sinks or floats based on what density?
A
when a solid object is placed into a liquid: object will FLOAT if it has a density less than that of a liquid object will SINK if it has a density greater than that of the liquid Temp decreases, density increases (inversely proportional)
17
Q
Covalent bonds and EN difference
A
18
Q
Dipole and EN
A
19
Q
Solvation Layer of hydrophillic & hydrophobic molecules
Order increases or decreases
A
-
Hydrophilic molecules in water: disorder increases (+S)
- water molecules form hydrogen bonds with hydrophilic groups
-
Hydrophobic molecules in water: order increases (-S)
- water molecues form a rigid hydrogen bond network around aggregated hydropobic groups
20
Q
Hydrogen Bonds donors & acceptors
A
- molecules with polar bonds that promotes dipole interactions with water are hydrophilic, whereas those with mostly nonpolar bonds lack attractive dipolar interactions with water are hydrophobic
- H atoms covalently bonded with N, O, or F atoms form POLAR BONDS that yeild dipoles, which exhibit signifcant noncovalent dipole-dipole attractions known as hydrogen bonding
21
Q
Sigma & Pi bonds
A
- Sigma Bonds: lower in energy, more stable, and have a greater dissociation energy than Pi bonds
-
Pi bonds: are weaker than sigma bonds
- higher energy state & are not as stable as sigma bonds, therefore require less energy to be broken than sigma (smaller dissociation energy)
-
Triple stronger > double> single
- b/c double and triple bonds are composed of both sigma and pi bonds
- double and triple bonds are composed of both sigma and PI bonds and are therefore stronger overall than a single bond
22
Q
Double and Triple Bonds
Example and how it related to bond dissociation energy
Relate it to rotation
A
-
single bonds are the longest: lowest bond dissociation energy & strength
- inc bond length
- free rotation (low bond rigidity)
-
double bond is medium
- 1 pi bond
- no rotation
-
triple bonds are the shortest: highest bond dissociation energy & strength
- 2 pi bonds
- dec bond length
- no rotation (inc bond ridigity)
23
Q
Intermediate Species
A
- When a reaction occurs in a stepwise sequence, the species formed as products in earlier steps and then subsequently consumed as reactants in later steps are intermediate species.
- Intermediates do not appear in the overall net reaction.
24
Q
Lewis Bases
A
- are electron pain donors
- stronger lewis bases can displace weaker lewis bases as ligands within a coordination complex
- Lone pair electrons on atoms with a lower EN tend to be stronger lewis bases than those with a higher EN
- charged atoms with lone pair electrons tend to be stronger lewis bases than comparable uncharged atoms
25
Example of STRONG LEWIS ACID and how it relates to radii
* metal cations with a **smaller ionic radius & a higher positive charge** such as *_Al+3 and Ti+4_* are **stronger lewis acids** than metal cations with a *larger ionic radius & a lower charge such as Ca+2 & Mg+2*
26
Percent Yield:
* **is a description of the efficiency of a reaction**
* It is the ratio of the mass obtained from an experiment (the actual yeild) to the calculated mass (theoretical yield) expressed as a percentage
* *actual yeild/theoretical \* 100%= % yield*
* **Example: table salt (NaCl) & AgNO3, how will the % yield be affected if some of the solid precipitate is lost when it is filtered from solution:** (solid AgCl is produced)
* If some of the solid AgCl is lost during filtration to isolate the product, the actual mass of the product obtained will be less than the calculated theoretical mass.
* **percent yield would be less than 100% becuase the mass of the isolated product is less than the calculated**
27
London Dispersion Forces:
* **also called induced dipole–induced dipole interactions (Van Der Waals)**
* **weakest of the noncovalent van der Waals forces**
* molecules exhibit weak mutual attraction
* these interactions occur between two atoms or molecules that are near enough to each other for distortions in the electron cloud to induce weak instantaneous dipoles.
* **LDF tend to be more pronounced in larger molecules with a larger, more polarizable electron cloud**
28
Dipole-Dipole Interactions
* are a **type of noncovalent interaction** that occurs between neighboring *polar molecules with polar bonds and a net dipole moment*
* The partial charges within nearby dipoles experience a mutual attraction and align in such a way that partial negative charges orient toward partial positive charges
29
Aufbau principle
* **low energy orbitals are filled first (4s fills before 3d),**
* **s** is LOWER ENERGY THAN **d**
* __**_The (n + l) rule_** can be used to rank subshells by increasing energy. This rule states that t*he lower the sum of the values of the first and second quantum numbers (n + l), the lower the energy of the subshell.*
* This is a helpful rule to remember for Test Day. If two subshells possess the same (n + l) value, the **subshell with the lower n value has a lower energy and will fill with electrons first.**
* For cations, to remove electrons, **you remove from the highest of energy subshells (highest n)**, if mulitple subshells are tied for the highest n value, then electrons are removed from subshells with the highest (L) value among these
30
Which of the following statements best explains why hemoglobin is red when it binds to O2?
* **O2 interacts with iron's d orbitals**
* **The nature of the ligands in a coordination sphere causes the metal's d orbitals to have different energies.**
* This energy difference determines the wavelength of light absorbed by the bonds.
* The wavelength of light that is reflected is often in the visible region of the electromagnetic spectrum.
31
**Coordinate Covalent Bonds**
* **are a special bond between a central atom and its ligands**
* The number of coordinate bonds indicates the *_coordination number_ and the ligands provide bonding electrons*, which interact with the *metal's d orbitals to form the coordinate covalent bond*
* **The metal ion Fe2+ is positively charged whereas the donor atoms (nitrogen) are neutral,** resulting in a net charge of +2 for the complex. Unlike ionic bonds, the donor atoms do not give electrons to the positively charged metal. Instead counterions often surround the complex in solution to balance the metal's positive charge.
32
Complex ion
* consists of a central atom, typically a metal ion, and its surrounding ligands, which form coordinate bonds to the metal
* The number of coordinate bonds to the central atom is the coordination number
33
Hybridization
* hybrid orbitals are formed by combining the atomic orbitals of an atom
* the hybridication of an atom is determined by counting the electron domains (sigma bond & lone pairs) & assigning a hybrid name
* The sum of superscripts must euqal the number of electron domains
34
Electromagnetic spectrum
* electromagnetic radiation at a certain wavelenght has photons of a particular energy
* **Energy is inversely proportional to wavelength**
35
Fluorescence
* When some of the energy is lost as heat and the rest is lost as light, the **emitted photon will have less energy than the absorbed photon, and therefore will have a longer wavelength.** *_This process is known as fluorescence._*
36
Excitation & Emission
37
Polar covalent & non-polar covalent bonds
Relate that to EN
* Covalent bonds involve electron sharing, and are considered **polar if the sharing is unequal but non-polar if the sharing is equal**
* **Polar Covalent:** atoms with an EN difference between 0.5 & 1.7 do not fully transfer electrons but instead share electrons unequally
* atom with higher EN gains the partial NEG charge and exerts a stonger pull on electrons
* **Nonpolar covalent:** forms between atoms with a difference in EN no greater than 0.5.
* frequently form between two atoms of the same type, such as C-C, or C-H
38
Electronegativity Numbers
39
which has the longest & weakest bond with the lowest bond dissociation energy?
**C-H**
**C//O**
**O-H**
**C-C**
* C-C BOND!!
40
Dipole moment and symmetry
* Molecules have a net dipole moment (separation of charge) when the individual dipoles within it do not cancel each other. **Symmetrical molecules typically do not have net dipole moments.** Protonation and deprotonation can change whether a molecule is symmetrical and can induce or remove a dipole moment.
**_Example:_**
* ph \< 1.25 is symmetrical so the dipoles cancel, giving no net dipole moment. Similarly, while the molecule is fully deprotonated at a pH \> 4.14, it is still symmetrical and has no net dipole moment in this form either
* **When the pH is between 1.25 and 4.14, oxalate is half protonated and is no longer symmetrical. In this form, it has a dipole moment.**
41
Example of a balanced Equation
42
Formal Charge
* is the charge assigned to an atom based on an accounting method, which assumes that the bonding electrons between two atoms are *_shared equally._*
* **FC= group valence - # of nonbonding electrons - 1/2 # of bonding electrons**
43
Valence Electrons
44
Double replacement and decompostion example
45
**Balancing redox reactions**
46
**Example of sigma/Pi bond of methionine**
**Which is true?**
* The pi bond in C=O has a smaller dissociation energy than the sigma bond.
* The C=O bond is stronger than the C–O bond.
* The pi bond in C=O is more stable than the sigma bond.
* **The pi bond in C=O has a smaller dissociation energy than the sigma bond**
* pi bonds exist in a **higher energy state** and **are not as stable as sigma bonds.** As a result, they require less energy to be broken than sigma bonds (ie, they have a **smaller dissociation energy)**
* **The C=O bond is stronger than the C–O bond**
* ****(Although individual pi bonds are weaker than sigma bonds, a **double bond is composed of both a sigma and a pi bond, and therefore is stronger than a single bond)**
* *3 is not TRUE, because Pi bond in C//O is higher in energy than the sigma bond, it is LESS STABLE*
47
Resonance
* Atoms never move, **only electrons.**
* All resonance structures must have the **same total number of valence electrons.**
* The **octet rule** must be obeyed fo**r first- and second-row elements.**
* Only electrons in pi bonds or lone pairs can move, **not electrons in sigma bonds.**
* **Shifting of electrons** should generally only be to **adjacent atoms** when going from one resonance structure to another.
* The **overall charge of the molecule must not change**; however, the formal charge of the constituent atoms can change.
48
Hybridization Example
49
**Valence shell electron pair repulsion (VSEPR)**
**Geometry of a molecule**
* **The geometry of a molecule can be determined according to valence shell electron pair repulsion (VSEPR) theory**
* *which states that the distribution of bonds and lone pairs will minimize repulsion between electrons.*
* **To find the geometry of a molecule,** count the number of electron domains and determine the electron group geometry, keeping in mind that lone pairs have greater repulsive forces than bonding pairs
* **if all electrons are from bonds**, the molecule's geometry will be the same as its electron froup geometry
* **if lone pairs are present**, the molecule's geometry will be different from electron group geometry
50
Constitutional isomers
51
Equilibrium Constant
Why is water ommited from the Keq expression?
* Because the **concentration** of the **solvent is much greater** than the concentration of all other species in the solution, **the solvent is relatively constant and is omitted from the Keq expression.**
52
Characteristics of IDEAL GAS:
53
**_Celsius to Kelvin Conversion:_**
54
**Average Kinetic Energy**
**Ideal gas**
**Boltzman constant**
* The *average kinetic energy* of a single ideal gas molecule is *directly proportional to temperature* via the **_Boltzmann constant,_** which can be expressed **_per mole of gas_** as the ideal gas constant R.
* According to the ideal gas law, PV is directly proportional to nT, and a graph of PV vs. nT results in a straight line with a slope of R.
55
STATE FUNCTION
56
**_What is boiling point influenced by?_**
* boiling point is influenced by **hydrogen bonding**
* *boiling point decreases as the hydrogen bonding decreases*
57
**Boiling point of related organic compounds**
58
**_Intermolecular Forces_**
* Structures with a **lesser** extent of **intermolecular forces** have **lower boiling points** because less energy is required to overcome the attractions between molecules during the **phase change** from a liquid to a gas state.
* When comparing similar structures, the boiling point tends to decrease as the number of hydrogen-bonding groups decreases due to fewer hydrogen bonding interactions.
59
Polar bonds
60
*Prediction of reaction shifts by comparing the reaction quotient Q to the equilibrium constant Keq*
61
What does it mean for something to be elastic or nonelastic?
62
Pressure measurement conversions
* Pressure measurements may be stated in a variety of units, including pascals (Pa), kilopascals (kPa), atmospheres (atm), millimeters of mercury (mmHg), and torr.
* *These units can be converted from one to another by using appropriate conversion factors, given that 1 atm = 760 mmHg = 760 torr = 101,325 Pa = 101.325 kPa.*
63
reaction rate
activaation energy & kinetic energy
* Molecules must collide with sufficient energy, known as the **activation energy,** for a reaction to occur.
* **_Increased temperature increases the kinetic energy of the molecules in a system, and therefore increases the rate of a reaction._**
64
Normal transition state: \_\_\_\_\_\_\_\_\_\_\_\_Ea
Stabilized transition state: \_\_\_\_\_\_\_\_\_\_\_Ea
* High Ea
* Low Ea
65
what does a catalyst do for the transition state?
* **stabilizes the transition state,** and therefore *increases the reaction rate without being consumed*
66
the amount of product formed in a reaction is dependent on what?
* **_on equilibrium constant of the reaction_**
* this constant is *not changed by a catalyst,* which only changes the rate at which equilirium is achieved
* same amount of products are produced when equilirbium constant is the same for both reactions
67
**1. Law of mass action states** that the rate of a reaction is *proportional* to \_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_
**2. Elementary reactions,** the reaction order of each species is equal to:
1. proporational to the **molar amount** (conc or Pa) of each reaction component **raised to the power of its reaction order**
2. equal to its **stoichiometric coefficient**
1. **for non-elementary reactions: the order must be determined empirically**
68
Henry's Law
* states that the amount of a **gas that dissolves in a liquid is proportional to the Pa of that gas**
69
A greater catalytic surface area corresponds to a **_lower/greater?_** reaction rate.
Catalysts provide a surface on which ___________ can form.
* **_GREATER!!_**
* **_example:_** a fine poweder provides more catalytic surfaces, & increases the speed of the reaction & the resulting power output
* *NOTE: smooth surfaces have less surface area than rough surfaces thus SLOWING reaction rates*
* *Catalysts provide a surface on which* **_stable transition states_** *can form*
70
**_In the standard state, Keq is related to spontaneity:_**
**_ΔG°:_**
* For a reaction to be spontaneous in the standard state, ΔG° must be negative.
* *Because ΔG° = −RT ln(Keq) and the natural logarithm of any number greater than 1 is positive, Keq must be greater than 1 for a standard state reaction to be spontaneous. Similarly, a Keq between 0 and 1 corresponds to nonspontaneous reactions, and a Keq equal to 1 corresponds to a reaction at equilibrium in the standard state. (detail at bottom)*
\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_
* + Ln (Keq)= **NEGATIVE** **_ΔG°; reaction is spontanous_**
* The natural logarithm (ln) of any number **greater than 1 is positive,** so **Keq must be greater than 1 (Keq\>1)** for the reaction to be *spontaneous* under standard state conditions.
* 0 Ln (Keq)= entire right side of the equation is equal to 0 and the system is at equilibrium in the standard state.
* **This occurs when Keq= 1 because ln(1) = 0**
* - Ln (Keq)= **positive _ΔG°; reaction is nonsponatous_**
* **__**The natural logarithm of any number between 0 and 1 is a negative number
71
**_Le Chatelier's Principle & Keq_**
72
**_Phase Diagram_**
73
**_Properties of Water:_**
*water is an efficient solvant due to its ability to interact with other oplar or charged compounds*
**_the polar nature of water is due?_**
* difference in the **_electronegativities_** of oxygen and hydrogen
* oxygen is more electronegative than hydrogen, the bonding electron pairs are more tightly drawn to oxygen, this allows oxygen to have partial negative charge and hydrogen a partial positive charge, allowing interations with other charged molecules
* along with a **_bent geometry_**
* **contributes to its polarity** by grouping charged positive charges at one end and negative on another end
* charged regions of a water molecule are sttracted to opposite charges on other polar compounds
* The **_small size_** of the water molecule allows efficient interactions with solutes, forming a hydration shell around and dissolving them.
* ***_Water has a high surface tension, but this characteristic does not aid in water's ability to act as a solvent._***
74
Surface tension
* **is a force induced at the interface between a liquid and a gas**.
* The molecules in the liquid interact with each other more strongly than they interact with molecules in the air, causing the surface of the liquid to behave as a thin film.
* **_Water has a high surface tension, but this characteristic does not aid in water's ability to act as a solvent._**
75
**_Raoult law and phase diagram_**
* **_Raoult law_** states that the addition of solute to a pure substance **lowers the freezing point and saturation vapor pressure of the resulting solution.** On addition of a solute, the phase boundary lines on a phase diagram should move accordingly.
76
**Freezing Point**
* The freezing point of a substance is **determined by the strength and extent of *_intermolecular forces_* acting between molecules.**
* These interactions are affected by the presence of solutes. As a colligative property, the extent of freezing point depression is a function of the amount of solute but not the solute's chemical properties.
* freezing point is the temperature at which solid & liquid phases of a substance are in equilirbium
* FP occurs when the kinetic energy of a molecule can no longer overcome the intermolecular forces binding to its nearby molecules
* **_example:_** if a molecule experiences strong intermolecular forces with surrounding molecules, the KE equired to overcome those forces is greater & the freezing point temp is correspondingly higher
77
INTRAmolecular forces
* forces **within** a molecule, & these forces do not impact interactions **betweeen** molecules
78
# Define a colligative property
**EXAMPLE?**
* **_Properties that are affected by the amount but not the identity (chemical properties) of the solute_**
* **_Example: FREEZING POINT_**
* ***__**the freezing point (a physical property) is not affected by the reactivity (a chemical property) of the solute.*
79
**_Hydrogen bonding:_**
* occurs when a partially positive hydrogen atom in one molecule is attracted to the partial negative charge on an electronegative atom in another molecule.
* **_Atoms involved in hydrogen bonding must have small atomic radii, and are essentially limited to fluorine, oxygen, and nitrogen._**
* large electronegative atoms (Br) cannot hydrogen bond
80
**_Vapor pressure_**
**_Boiling Point_**
* **Vapor pressure**: is the Pa of the gas that is released from the liquid phase at a any givne temperature
* **on phase diagram**: VP is shown as the boundry line between the gas & liquid phases & it increases with temperature
* *when temp is high enough for the vapor pressure to equal the ambient pressure, the liquid BOILS*
* **Boiling point** temperature is the temperature at which the **vapor pressure=ambient pressure.**
* The **addition of solute** *lowers the vapor pressure of a solution at all temperatures (Raoult Law),* and therefore *raises the temperature required for the vapor pressure to become equal ambient pressure and begin to boil (the boiling point).*
81
Irreversible Reactions
Reversible Reactions
* **Irreversible Reactions**
* *Low temperatures*
* under **kinetic control**
* _*most favorable product is the one whose pathway has the **LOWEST Ea***_
* fastest reaction will yield the most product
* **Reversible Reactions**
* *High temperature*
* *thermodynamic*
* the ***_most stable product_*** is favored **(most negative ΔG),** as it is least likely to undergo the reverse reaction.
82
**_Rate Law Types_**
83
Effect of temperature on Gibbs Free Energy
* **the sign of the change in ΔS for a process determines** whether the process becomes **more or less thermodynamically favorable** as the temperature increases.
* *_A negative ΔS causes ΔG to increase with temperature, whereas a positive ΔScauses ΔG to decrease with temperature._*
84
**_Arrhenius Equation_**
* describes the relationship between the **rate constant and the temperature and activation energy of a reaction.**
* Both the rate constant and the rate of reaction decay exponentially with **increasing activation energy and increase exponentially with increasing temperature**.
* Because catalysts lower the activation energy of a reaction, the **rate constant is greater for a catalyzed** reaction than for an uncatalyzed reaction
*
85
**_Equilirbrium Constant and Le Châtelier principle & Temperature_**
* Equilibrium is reached in a reaction when the rate at which products are formed is equal to the rate at which reactants are formed.
* **The equilibrium constant Keq** is defined as the ratio of the reactants and products when a reaction is in equilibrium, and **_only changes in temperature can alter Keq._**
* **_Le Châtelier principle_** states that a system responds to disturbances (temp, pressure, or component's conc) in equilibrium by shifting the reaction to counter the effect of the disturbance.
* changes in pressure & conc achieve equilibrium w/o altering the equilirbium constant Keq
86
Oxidation states
87
**_Osmotic Pressure Example:_**
**_Van Hoff's Factor_**
88
**Redox reaction balancing Example**
89
**_Hybridization_**
**_Trigonal Bipyramidal_**
* **Hybrid orbitals form when two or more atomic orbitals (s, p, d, f) combine into new orbitals.**
* *as the number of electron-dense areas increase, repulsion around the orbital increases*
* *Electrons distribute themselves into specific configurations that minimize this repulsion. These electron distributions act as frames in which bonding may occur.*
* **The number of hybrid orbitals** (electron-dense areas) around an atom dictates its **hybridization and electron geometry,**
* whereas only the **orientation of *_bonds_*** around the central atom determines the **molecular geometry (shape) of a molecule.**
* ****non-bonding electrons are ignored
* *Two compounds may have the same number of electron-dense areas but different molecular geometries if they differ in number of lone electron pairs.*
* *EXAMPLE:*
* dsp3 have 5 electron-dense areas, 1 from d, 1 from s, and 3 from p orbitals
* To minimize repulsion, three of the hybrid orbitals lie in the same plane, 120° apart, and the other two orbitals are 90° above and below the plane. **_Compounds with dsp3 hybridization and no lone pairs are trigonal bipyramidal._**
* if a compoud with dsp3 hybridization has either one or two lone pairs, it will have a see-saw or T-shaped geometry, respectively
90
See-Saw
91
**T-shaped geometry**
92
**_tetrahedral molecular geometry_**
* has 4 electron-dense areas around an sp3 hybridized central atom
93
Octahedral
* no lone pairs
* has 6 electron-dense areas around a d2sp3 hybridized central atom
94
square planar
* (two lone pairs) molecular geometries have six electron-dense areas around a d2sp3 hybridized central atom.
95
Kw, pH
96
Precipitate formation reaction
* reactions producing an insoluble solid (ppt) crashing from the solution
97
chelate formation reaction
* a metal cation & a ligand react to form one or more rings via a pincer-like coordinate bonding arrangement
98
Oxidation-reduction (redox) reaction
* the oxidation of some atoms change during the conversion of the reactant to the products
* example: Shiny nickel metal granules added to orange liquid bromine mixed in alcohol produced a combination reaction yielding a blue-green solution of nickel(II) bromide. In the reaction between Ni and Br2: **_a nickel undergoes an oxidation-reduction with bromine_**
* **_explained:_** *The given reaction between Ni and Br2 is a combination reaction in which two reactants form a single ionic product, NiBr2. The assigned oxidation states for the elements in this reaction show that nickel goes from an oxidation state of zero in its elemental metal form to a state of +2 in the product (an oxidation), and bromine goes from an oxidation state of zero to a state of −1 (a reduction). Therefore, in addition to being a combination reaction, the reaction between Ni and Br2 is also an oxidation-reduction reaction.*
99
Effective nuclear charge
* outer electrons are shielded from the pull of protons by the inner electrons
* **electrons experience a decreased attractive force toward the nucleus (Zeff), so they orbit farther from the nucleus**
100
Atomic Mass Unit (amu)
is defined as one-twelfth of the mass contained in a carbon-12 atom (the average mass of 6 protons and 6 neutrons). On this basis, the amu provides a useful measure for assessing the masses of different atoms relative to the carbon-12 standard.
101
the molar mass concept
102
Diatomic gases and their bonds
103
order of orbital filling, in increasing energy
104
Ionic radii, Zeff, and isoelectronic series
* ionic radii tend to decrease in size across a period (row) (left to right) and increase moving down a group (column)
* Compared to the neutral atom of a given element, its **cation will be smaller but its anion will be larger.**
* *Losing electrons* to form a cation causes the remaining electrons to experience a **greater effective nuclear charge Zeff, pulling the electrons closer to the nucleus.**
* Conversely*, gaining electron*s to form an anion produces greater electronic repulsion and nuclear shielding **(lesser Zeff),** which pushes electrons farther from the nucleus.
* The ions S2−, Cl−, K+, and Ca2+ are isoelectronic (have the same number of electrons); however, because the number of protons is different in each ion, the electrons in each ion experience a different Zeff.
* Therefore, ***_in an isoelectronic series, ionic radii decrease as the atomic number increases._*** Because sulfur has the lowest atomic number (fewest number of protons) in this isoelectronic series, Zeff is smallest in S2−, making it the largest ion. Given that the cell channel preferentially transmits smaller ions, S2− is least likely to be transmitted.
105
Determining oxidation states
106
Formal charge
107
How does oxidation differ from formal charge
* Oxidation numbers differ from the formal charge in that oxidation numbers account for **redox changes** whereas formal charge **accounts for charge distribution among atoms.**
* However, the sum of all the oxidation states of every atom in the compound must equal the net charge of the compound formula.
108
Disporportionation reaction
* In a disproportionation reaction, the **same element (at a given oxidation state) undergoes both oxidation and reduction,** with some of the atoms being oxidized and other atoms of the *same element* being reduced. Comparing the oxidation number for each element in the reactants with the oxidation number for the same element in the products enables the identification of a disproportionation.
* oxidation numbers for elements oxidation increases, whereas those for elements undergoing reduction decrease the oxidation number
109
Electrical conductivity tendency
* **_Metals_** tend to be shiny, ductile, and good conductors of heat and electricity whereas **_nonmetals_** tend to be dull, brittle, and poor electrical conductors (making them good electrical insulators).
* **_Metalloids_** share some characteristics of metals and some characteristics of nonmetals. As a result, metalloids tend to be better electrical conductors than nonmetals but not as conductive as metals. This makes metalloids useful as semiconductors in some instances and as insulators in others.
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Representative elements:
* (main group elements) are the elements in groups 1-2 (1A-2A) & groups 13-18 (3A-8A) that are contained in the s-block and the p-block, repectively, of the periodic table
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first, second and 3rd ionization energry
* Ionization energy is the energy required to remove an electron from a ground-state gaseous atom or ion. The first, second, and third ionization energies describe the energy required to remove one, two, or three electrons, beginning with the electron of highest energy. **Electrons in the inner core require significantly more energy to remove than valence electrons.**
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order of orbital removal of electrons
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ionization energy and reactivity ratio
* **_inversly proportional_**
* The alkali metals are very reactive because of their **low ionization energy, large atomic radius, and small electronegativity**. *Cesium is considered the most reactive metal.*
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Hydrogen halides
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**Hydrogen halides**
Example: H-F is not strong acid, why?
the acid strength of hydrogen halides is ranked from the strongest to weakest as follows:
**H-I \> H-Br \> H-Cl \>\>H-F**
* The acid halides are all strong acids except for H−F. **_Fluorine is the most electronegative and has the smallest atomic radius of the halides, resulting in strong bond formation (H-F is not a strong acid)_**
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Electron affinity and electron affinity proportion
* Electron affinity is the measure of energy change when an electron is added to an atom in the gaseous state.
* **Electron-electron repulsions result in a decreased electron affinity.**
* example: even though oxygen is the most electronegative element in group 6A, it has the lowest electron affinity in the group becuase OXYGEN'S VALENCE ELECTRONS EXPERIENCE MORE ELECTRON-ELECTRON REPULSION THAN OTHER GROUP 6A ELEMENTS
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What is Gd+3 and N are classified as in the picture?
**_coordination chemistry_**
Gd+ is a lewis acid
N are lewis bases
* In *coordination chemistry,* metal ions such as Gd3+ can act as Lewis acids because they have empty valence orbitals that can readily accept electrons. Molecules or atoms with lone pairs of electrons act as Lewis bases and coordinate to the metal ion, forming a complex ion; these electron-donating groups are referred to as ligands. In this example, the ligand molecules are amide nitrogen residues and water molecules.
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During ionization, which electrons electron are removed from the valence electron shell?
the ones with the greatest principal quantum number first becuase these electrons are the farthest from the nucleus and thus are the least tightly bound
Example: Gd+3 will lose its first two from the 6S subshell followed by one electron from the 5d subshell to produce Gd+3, which has the configuration of [Xe]4f
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Heisenberg uncertainty principle
* The Heisenberg uncertainty principle states that the **position and momentum of a particle cannot both be precisely known at the same time.** The uncertainty of the position and the uncertainty momentum are inversely proportional.
* "electrons in coordination bonds can be described only as probability distributions" is a consequence of this principle
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**_Types of magnetism_**
* Hund's rule states that orbital filling maximizes the number of unpaired electrons.
* **Paramagnetic atoms** and molecules have unpaired electrons that align parallel to an applied magnetic field; paramagnetic species are weakly attracted to magnets.
* In contrast, **diamagnetic atoms** and molecules have no unpaired electrons; these species are repelled by magnets.
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Dilution factor calculation
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