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U1.2 Test Chemistry
Unit 1.2: Electron Arrangement and Periodic Law Chemistry Review
| Question | Answer |
|---|---|
| An unknown wave has a frequency of 3.0x10^18 Hz. What is the wavelength in meters? | c = f*lang lang = c/f (3.00x10^8 m/s)/(3.0x10^18 s^-1) = 1.0x10^-10m |
| Rank the types of waves in decreasing energy, x-rays, AM radio, UV rays, gamma rays, FM radio, infrared rays, blue light, purple light, yellow light, green light, red light, and orange light. | gamma, x-ray, uv, purple, blue, green, yellow, orange, red, infrared, FM, AM |
| The red light from a helium-neon laser has a wavelength of 633 nm. What is the energy of the photon? Express your answer with units and appropriate number of sig figs. | E = hc/lang <-- must be in m 633nm = 6.33x10^-7m ((6.63x10^-34 J*s)(3.00x10^8 m/s))/6.33x10^-7m = 3.14x10^-19 J |
| Given the following electron configuration, describe what each letter and number means: 1s^2 | 1 = principle energy level s = sublevel 2 = number of e- (electrons) in that sublevel |
| Identify the element that has the following electron configuration: 1s^2 2s^2 2p^6 3s^2 | Magnesium |
| Identify the element that has the following electron configuration: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^2 | Titanium |
| Identify the element that has the following electron configuration: [Xe] 6s^1 | Cesium |
| Explain why atomic radius of an atom gets smaller as you move from left to right across a period (possible terms to use: shielding, coulombic forces of attraction and nuclear charge) | Nuclear charge increases --> pulling e- (electrons) in closer |
| Explain why ionization energy decreases as you move from top to bottom down a group (possible terms to use: shielding, coulombic forces of attraction and nuclear charge) | More shielding (inner e- blocking + nucleus from outer e-) shielding = large atoms blocking = makes it easier to remove outer e- ---> meaning to lower i.e. |
| Explain what is meant by "Coulombic forces of attraction" | + and - attract each other, but the greater distance between the 2 ---> the weaker the attraction |
| Choose the element that exhibits the characteristics below: B, Li, F. Forms a 1+ ion | Li |
| Choose the element that exhibits the characteristics below: B, Li, F. Smallest atomic radius | F |
| Choose the element that exhibits the characteristics below: B, Li, F. Most reactive nonmetal | F |
| Choose the element that exhibits the characteristics below: B, Li, F. Electron configuration ends in 2p^1 | B |
| Choose the element that exhibits the characteristics below: B, Li, F. Highest electronegativity | F |
| Choose the element that exhibits the characteristics below: B, Li, F. Forms a 1- ion | F |
| Choose the element that exhibits the characteristics below: B, Li, F. Metal | Li |
| Choose the element that exhibits the characteristics below: B, Li, F. Forms 3+ ions | B |
| Choose the element that exhibits the characteristics below: B, Li, F. Lowest ionization energy | Li |
| Choose the element that exhibits the characteristics below: B, Li, F. Nonmetal | F |
| Choose the element that exhibits the characteristics below: B, Li, F. Most reactive metal | Li |
| Choose the element that exhibits the characteristics below: B, Li, F. Has 2 empty orbitals | B |
| Choose the element that exhibits the characteristics below: B, Li, F. Properties of metals and nonmetals | B |
| Choose the element that exhibits the characteristics below: B, Li, F. 7 valence electrons | F |
| Choose the element that exhibits the characteristics below: B, Li, F. Soft, very malleable metal | Li |
| Choose the element that exhibits the characteristics below: B, Li, F. Electron configuration ends in 2p^5 | F |
| Write full electron configuration of each element: Nickel (Ni) | 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^8 |
| Write full electron configuration of each element: Sulfur (S) | 1s^2 2s^2 2p^6 3s^2 3p^4 |
| Write noble gas electron configuration of each element: Nickel (Ni) | [Ar] 4s^2 3d^8 |
| Write noble gas electron configuration of each element: Sulfur (S) | [Ne] 3s^2 3p^4 |
| Write number of valence electrons of each element: Nickel (Ni) | 2 |
| Write number of valence electrons of each element: Sulfur (S) | 6 |
| Write number of core electrons of each element: Nickel (Ni) | 26 |
| Write number of core electrons of each element: Sulfur (S) | 10 |
| Write the charge of each element: Nickel (Ni) | +2 |
| Write the charge of each element: Sulfur (S) | -2 |
| The Bohr Model of the atom was an attempt to explain hydrogen's | atomic emission spectrum |
| A bright-line spectrum of an atom is caused by the energy released when electrons: | fall to a lower energy level |
| The lowest energy state of an electron is called the: | ground state |
| What is the max number of orbitals in the d sublevel? | 5 |
| The major difference between a 1s orbital and a 2s orbital is that: | the 2s orbital is at a higher energy level |
| How many unpaired electrons are there in a selenium atom (atomic number 34)? | 2 |
| Identify the element with the following electron configuration: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^3 | Arsenic |
| If three electrons are available to fill 3 empty 2p atomic orbitals, how will the electrons be distributed in the three orbitals? | One electron in each orbital |
| The group of soft, silvery active metals, all of which have 1 electron in an s orbital, is known as the: | Alkali Metals |
| What is the chief characteristics of the noble gases? | Very low reactivity |
| Each period in the periodic table corresponds to: | A principle energy level |
| What is true of the electron configuration of the noble gases? | The outermost s and p sublevels are filled |
| What orbital is filled when iodine gains an electron to become a negative ion? | 5p |
| Valence electrons are those | in the highest energy level |
| How many valence electrons does an atom of any halogen have? | 7 |
| The number of valence electrons in Group 2 elements is: | 2 |
| When an aluminum atom loses its valence electrons, the charge on the resulting ion is | 3+ |
| To attain a noble gas configuration a sulfur atom must: | gain 2 electrons |
| How does atomic radius change from left to right across a period in the periodic table | it tends to decrease |
| How does atomic radius change down a group in the periodic table? | it tends to increase |
| Heisenberg's uncertainty principle states that | it is impossible to know the position and velocity of an electron simultaneously |
| Match with correct atomic history: Discovered the neutron | Rutherford |
| Match with correct atomic history: Modeled the orbitals as electron cloud with different shapes and sizes | Schrodinger |
| Match with correct atomic history: Confirmed the existence of neutrons | Chadwick |
| Match with correct atomic history: Determined the charge on an electron | Thomson |
| Match with correct atomic history: Greek philosopher who named the smallest particle of matter "atomos" | Democritus |
| Match with correct atomic history: Believed that all atoms of an element were identical | Dalton |
| Match with correct atomic history: The electron could circle the nucleus only with in allowed paths (orbits) | Bohr |
| Match with correct term: Arrangement of electrons around an atoms nucleus | Electron Configuration |
| Match with correct term: No more than two electrons can occupy an atomic orbital and these two electrons must have opposite spin | Pauli Exclusion Principle |
| Match with correct term: Region of high probability of finding an electron | Orbital |
| Match with correct term: Tendency of electrons to enter orbitals of lowest energy first | Aufbau Principle |
| Match with correct term: When electrons occupy orbitals of equal energy, one electron enters each orbital until all orbitals contain one electron with parallel spins | Hund's rule |
| Compare and contrast metals, nonmetals, and metalloids | metals: high melting points, ductile, malleable, forms cation, shiny luster, good conductors nonmetals: form anions, brittle, dull, low melting point, and non-conductive metalloids: metal and nonmetal properties (semiconductor industry) |
| According to Bohr Model of hydrogen atom, how is hydrogen's emission spectrum produced? | smaller fall --> (red) less energy larger fall --> (violet) more energy electrons get hit with a proton causing it to jump up energy levels, becoming excited. electrons can't stay excited, so it falls back to ground state, releasing E as light. |
| Describe Mendeleev's contribution to the periodic table, why did he leave gaps in his table? What were the major limitations of his periodic table? | 1st periodic table --> gaps because not all elements were known yet. Didn't arrange by atomic number |
| How does Moseley's periodic arrangement of the elements differ from Mendeleev's? | Moseley used atomic number while mendeleev arranged by atomic mass |
| Match the correct term: When electrons in inner energy levels block the attraction of the nucleus for the valence electrons | Shielding |
| Match the correct term: The energy required to remove one electron from a neutral atom of an element | Ionization Energy |
| Match the correct term: A negative ion formed by the gain of electrons | Anion |
| Match the correct term: A positive ion formed by the loss of electrons | Cation |
| Match the correct term: The tendency for an atom to gain an electron | Electronegativity |
| Match the correct term: The strength of the nucleus to hold (attract) the outer electrons due to an increase of protons in the nucleus | Nuclear Charge |
| Atomic radii generally ________________as you move from left to right in a given period because there is an increase in the nuclear charge while the number of inner electrons, and hence the shielding effect, remains constant. | Decrease |
| Atomic radii generally _________________ within a given group because the outer electrons are farther from the nucleus as you go down the group. | Increase |
| The attractive force of the increased nuclear charge is unable to overcome the effect of the greater distance, which acts in opposition. | memorize this |
| Ionization energy generally ____ as you move from left to right across a period. | increases |
| Ionization energy ____________ as you move down a group. | decreases |
| Electronegativity generally ___________________ as you move from left to right across a period. | increases |
| Electronegativity ____________ as you move down a group. | decreases |
| Why do positive ions have smaller ionic radii than neutral atoms? Example: Why is the lithium atom larger than the lithium ion? | Loss of electrons Greater nuclear-charge compared to electrons --> so pulls in tighter |
| Why do negative ions have larger ionic radii than neutral atoms? Example: Why is the chloride ion larger than the chlorine atom? | Gain of electrons Cl +17 (17e-) --> Cl +17 (18e-) |
| Identify the scientist and atomic theory from atomic models: Quantum model | Schrodinger 3-D |
| Identify the scientist and atomic theory from atomic models: Planetary Model | Bohr |
| Identify the scientist and atomic theory from atomic models: Plum pudding model | Thomson |
Created by:
Sabrina_thecool
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