Periodic Table
Covalent-nonpolar bond Covalent-polar bond Ionic bond
A __________________ bond is a chemical bond where the electrons in a bond are shared equally. A ___________________ bond is a chemical bond where the electrons in a bond are shared unequally. A ___________________ bond is a chemical bond where the electrons in a bond are not shared, one atom takes both electrons from the other atom it is bonded to.
Greater Less
An atom with a ______________ nuclear charge increases ionization energy, while an atom with ________ nuclear charge decreases ionization energy.
Answer: B. Antimony (Sb) is on the right side of the periodic table, but not far right enough to be a nonmetal, (D.). It certainly does not lie far right enough to fall in Group VIIA (Group 17), which would classify it as a halogen, (C.). While sources have rarely classified antimony as a metal, (A.), it is usually classified as a metalloid, (B.).
Antimony is used in some antiparasitic medications - specifically those targeting Leishmania donovani. What type of element is antimony? A. Metal B. Metalloid C. Halogen D. Nonmetal
1. Any elements that are considered metals 2. Any elements that are considered nonmetals 3. Any elements that are considered metalloids
Based on their location in the periodic table, identify a few elements that are likely to possess the following properties: 1. Luster 2. Poor conductivity and electricity 3. Good conductivity but brittle
Answer: A. As one moves from top to bottom in a group (column), extra electron shells accumulate, despite the fact that the valence configurations remain identical. These extra electron shells provide shielding between the positive nucleus and the outermost electrons, decreasing electrostatic attraction and increasing atomic radius. Because carbon and silicon are in the same group, and silicon is farther down in the group, silicon will have a larger atomic radius because of its extra electron shell.
Carbon and silicon are the basis of biological life and synthetic computing, respectively. While these elements share many chemical properties, which of the following best describes a difference between the two elements? A. Carbon has a smaller atomic radius than silicon B. Silicon has a smaller atomic radius than carbon C. Carbon has fewer valence electrons than silicon D. Silicon has fewer valence electrons than carbon
Water Neutral state Less dense Reactive Valence electron Z(eff) Effective Nuclear Charge Large Low Low Low One
Characteristics of alkali metals: -Highly reactive to ________ -Not found in _________ _________ in nature -________ __________ than other metals -Extremely _____________, easily able to lose ___________ ___________ to form a univalent cation. -They have a very low _____________ -___________ atomic radii, _______ electronegativities, ________ ionization energies, _______ electron affinity. -Have ________ valence electron
Water Neutral State Reactive Valence electrons Higher Smaller More Greater Two
Characteristics of alkaline earth metals: -Highly reactive to ________ -Not found in __________ _________ in nature -Extremely ___________, easily able to lose ____________ _____________ to form a divalent cation. -Z(eff) charge is slightly ___________ than alkali metals -Slightly ___________ atomic radii than alkali metals -Slightly ____________ electronegative than alkali metals -Slightly ____________ ionization energy and electron affinity than alkali metals. -Have ___________ valence electrons
Nonmetals Metals and metalloids Atomic radii electronegative ionization energy electron affinity Six
Characteristics of chalcogens: -____________ are crucial for biological functions -___________ and __________________ are genrally toxic -Small _____________ ___________, highly __________________, great _________________ ______________ and _________________ ________________ -Have ________ valence electrons
reactive neutral form Smaller More Greater Seven
Characteristics of halogens: -Very _____________ nonmetals, they don't appear in ___________ __________ in nature. -_____________ atomic radii than chalcogens -____________ electronegative than chalcogens -______________ ionization energy and electron affinity than chalcogens -Has ___________ valence electrons
Good Metals luster brittle
Characteristics of metalloids: -________ conductivity of electricity, not as great as ___________ -May have a metallic __________ -___________, not very malleable.
solid malleable ductile good
Characteristics of metals: -Typically ___________ at room temperature -_____________, able to alter its shape -____________, can be drawn into wires -__________ conductors of heat and electricity
Ionization energy Electron affinity Unreactive Bond Eight
Characteristics of noble gases: -High ________________ __________________ -Low/nonexistent ________________ ______________ -Very ______________, not very willing to ___________ -Has __________ valence electrons
states of matter brittle poor
Characteristics of nonmetals: -Can be found in a variety of different ___________ of _____________ at room temperature - ___________, they are not very malleable when solid -_________ conductors of electricity
low low low hard high oxidation states d-orbitals
Characteristics of transition metals: -_______ electron affinities, ________ ionization energies, ________ electronegativities. -Very ______, __________ melting and boiling points -Are able to exist in multiple different _____________ __________, due to having ______________ in their valence shell.
Ag - M Al - M K - M P - NM Pb - M Li - M Pu - M U - M Cu - M As - MO Zn - M B - MO Si - MO Np - M O - NM He - NM
Classify the following elements as metals (M), nonmetals (NM), or metalloids (MO): -Ag -Al -K -P -Pb -Li -Pu -U -Cu -As -Zn -B -Si -Np -O -He
greater greater
Due to the removal of the first electron, the attractive force on the valence electron will be much _____________ than previously, causing the valence to be drawn closer to the nucleus. By getting closer to the nucleus, the attractive force exerted on the valence electron will be ______________ than before first ionization. This results in a massive increase in ionization energy for removing another electron.
Left Right Down
Effective nuclear charge (Z(eff)) increases from _______ to _____________ on the periodic table. Principle quantum number increases by 1 the further __________ you go on the periodic table.
Representative elements Nonrepresentative elements
Elements in Groups 1-2 and 13-18 are known as ______________________ elements. Elements in Groups 3-12 and the lanthanide and actinide elements are known as ___________________ elements.
1. Group 1 and 2 2. Group 6 and 16 3. Groups 1-15 4. All groups (most notably Groups 3 - 12) 5. Almost all groups (most notably Groups 14 - 17) 6. Group 18
For each of the properties listed below, write down the groups of the periodic table that exhibit those properties. 1. High reactivity to water 2. Six valence electrons 3. Contain at least one metal 4. Multiple oxidation states 5. Negative oxidation states 6. Possess a full octet in the neutral state
Group 1 Group 2 Group 16 Group 17 Group 18
Group ____ on the periodic table consists of alkali metals. Group ____ on the periodic table consists of alkaline earth metals. Group ____ on the periodic table consists of chalcogens. Group ____ on the periodic table consists of halogens. Group ____ on the periodic table consists of noble gases.
Negligible
Helium, neon, and argon are the three elements that have _________________ electronegativities.
3 same extra opposite repelled electron shielding decrease
How Elements in Group 16 defy ionization trend: Group 15 elements have ____ electrons in the p-subshell, with each of them in their own orbital spinning in the ___________ direction. Group 16 elements has a ________ electron added to one of those orbitals, but spinning in the _______________ direction. This causes the electron being added to be ______________ by the electron that was already present in that orbital, causing more _______________ ____________, causing a ______________ in effective nuclear charge, resulting in the decrease in ionization energy.
p-orbital higher p-orbital farther atomic radius p-orbital repelled electron shielding lower
How Group 13 elements defy ionization trend: Group 13 elements has an electron that is added to the ____-orbital, which is ____________ in energy than the s-orbital. Electrons in the _____-orbital are _____________ from the nucleus than the s-orbital, which increases ___________ ______________, which decreases ionization energy. Furthermore, the placement of of the electrons in the ____-orbital will be _______________ by the electrons in the s-orbital along with the other electrons in the inner shells. This leads to an increase in _____________ _____________, which leads to a ____________ effective nuclear charge, which decreases ionization energy
Same Repelled Increase No
How Group 15 elements defy electron affinity trend: Group 15's p-orbitals all have an electron in each of them spinning in the ___________ direction. The addition of an extra electron into one of these orbitals will cause it to be _________________ by the electron that is already present. This leads to a ________________ in electron shielding, which leads to _______ attraction between the valence electron and nucleus. This results in nitrogen not having any electron affinity.
Higher 2p6 3s1 Shielded 0 No
How Group 18 elements defy electron affinity trend: When we add an electron to a Group 18 element, it will cause a valence of a ____________ energy level to be formed. In the case of neon, we go from a _______ subshell to a _____ subshell. This addition will cause the newly added electron to be ___________ by all the inner electrons. With all of this considered, Group 18 elements will have an effective nuclear charge of ___, meaning that the added electron will have ____ attraction to the nucleus, meaning it does not have an affinity for an electron.
p-subshell shielded Electron shielding Energy
How Group 2 elements defy electron affinity trend: When adding an electron to a Group 2 element, the electron will be added into the ________-subshell instead of the s-subshell. Because of the new subshell it is placed in, it will be _______________ by both the 1s electrons, but also the 2s electrons. The increase in ______________ ____________ will cause the added electron to not exhibit attraction to the nucleus, meaning it would take ___________ to add the electron. This results in it not having an electron affinity
Answer: D. The question is simple if one recalls that periods refers to the rows in the periodic table, while groups or families refer to the columns. Within the same period, an additional valence electron is added with each step towards the right side of the table.
How many valence electrons are present in elements in the third period? A. 2 B. 3 C. The number decreases as the atomic number increases D. The number increases as the atomic number increases
Electronegative
In a polar covalent bond, the atom that has a greater attraction of the shared electrons is the atom that is more _______________________.
F- K
In each of the following pairs, which has the larger radius? -F or F- -K or K+
Group 2 Group 15 Group 18 decreasing
In the periodic trend of electron affinity, elements in Group ______, Group _____, and Group _____ defy this trend by _________________ in electron affinity from the preceding element.
Group 13 Group 16 Decreasing
In the periodic trend of ionization energy, elements in Group ______ and Group _____ defy this trend by _________________ in ionization energy from the preceding element.
Answer: C. Ionization energy increases from left to right, so the first ionization energy of lithium is lower than that of beryllium. Second ionization energy is always larger than first ionization energy, so beryllium's second ionization energy should be the highest value. This is because removing an additional electron from Be+ requires one to overcome a significantly larger electrostatic force.
Ionization energy contributes to an atom's chemical reactivity. Which of the following shows an accurate ranking of ionization energies from lowest to highest? A. first ionization energy of Be < second ionization energy of Be < first ionization energy of Li B. first ionization energy of Be < first ionization energy of Li < second ionization energy of Be C. first ionization energy of Li < first ionization energy of Be < second ionization energy of Be D. first ionization energy of Li < second ionization energy of Be < first ionization energy of Be
Endothermic process Exothermic process
Ionization energy is calculated from an ___________________ process, since the removal of electrons requires energy to execute. Electron affinity is calculated from an ___________________ process, since the addition of an electron dispels energy in the form of heat.
kJ/mol
Ionization energy is measured in ____________.
Answer: B. The periodic table is organized into periods (rows) and groups (columns). Groups (columns) are particularly significant because they represent sets of elements with the same valence electron configuration, which in turn will dictate many of the chemical properties of those elements. Although (A.) is true, the fact that both ions are positively charged does not explain the similarity in chemical properties; most metals produce positively charged ions. (C.) is not true because lithium and sodium do have relatively low atomic weights, so do several other elements that do not share the same properties, eliminating (D.).
Lithium and sodium have similar chemical properties. For example, both can form ionic bonds with chloride. Which of the following best explains this similarity? A. Both lithium and sodium ions are positively charged B. Lithium and sodium are in the same group of the periodic table C. Lithium and sodium are in the same period of the periodic table D. Both lithium and sodium have low atomic weights
Atomic number
Mendeleev's table was arranged by atomic weight, but the modern periodic table is arranged by ____________ ___________.
Answer: C. All four descriptions of metal are true, but the most significant property that contributes to the ability of metals to conduct electricity is the fact that they have valence electrons that can move freely. Malleability, (A.), is the ability to shape a material with a hammer, which does not play a role in conducting electricity. The low electronegativity and high melting points of metal, (B.) and (D.), also do not play a major role in the conduction of electricity.
Metals are often used for making wires that conduct electricity. Which of the following properties of metals explains why? A. Metals are malleable B. Metals have low electronegativities C. Metals have valence electrons that can move freely D. Metals have high melting points
Easier Decreases Harder Increases
More electron shielding makes a valence electron less attracted to the nucleus, which makes it ____________ for it to be removed. Therefore, more electron shielding _______________ ionization energy. Less electron shielding makes a valence electron more attracted to the nucleus, which makes it ______________ to remove it. Therefore, less electron shielding ______________ ionization energy.
Answer: B. Electron affinity is related to several factors, including atomic size and filling of the valence shell. As atomic radius increases, the distance between the nucleus and the outermost electrons increases, thereby decreasing the attractive forces between protons and electrons. As a result, increased atomic radius will lead to lower electron affinity. Because atoms are in a low-energy state when their outermost valence electron shell is filled, atoms needing only one or two electrons to complete this shell will have high electron affinities. In this example, (B.) and (D.) need only one more electron to have a noble gas-like electron configuration; because (B.) is smaller, it will have the highest electron affinity.
Of the four atoms depicted here, which has the highest electron affinity?
1. O 2. Sb 3. Tl 4. Ne
Rank the following elements by decreasing electronegativity: -Antinomy (Sb) -Neon (Ne) -Oxygen (O) -Thallium (Tl)
1. C 2. Ge 3. Ca 4. K
Rank the following elements by decreasing first ionization energy: -Calcium (Ca) -Carbon (C) -Germanium (Ge) -Potassium (K)
1. Xe 2. Nb 3. Ta 4. Pr
Rank the following elements by increasing atomic radius: -Niobium (Nb) -Praseodymium (Pr) -Tantalum (Ta) -Xenon (Xe)
1. Ba 2. Y 3. Cu 4. S
Rank the following elements by increasing electron affinity: -Barium (Ba) -Copper (Cu) -Sulfur (S) -Yttrium (Y)
Greater Smaller
The _____________ the effective nuclear charge, the greater the ionization energy, the _______________ the effective nuclear charge, the smaller the ionization energy.
Atomic radius
The average distance between a nucleus and its outermost electron is known as the ___________ ___________.
Ionic radius
The average distance from the center of the nucleus to the edge of an electron cloud for a cation and anion is known as the ______________ ____________.
Increases Decreases
The closer the electron is to the nucleus, the greater the attractive force will be. Thus, less distance between the nucleus and valence electron ______________ ionization energy. The further the electron is to the nucleus, the weaker the attractive force will be. Thus, more distance between the nucleus and valence electron __________________ ionization energy.
Ionization energy
The energy required to remove an electron from one mole of gaseous atoms to produce one mole of gaseous ions is known as the ________________ __________________ of an atom. X + energy -----> X+ + e-
Increases Greater Decreased Increase
The further right you go on the periodic table, the nuclear charge of the elements ________________ by 1, which causes a ____________ attractive force of the valence electron to the nucleus, which also leads to a ______________ distance between the valence and nucleus, resulting in an _____________ in ionization energy.
Larger Smaller Smaller Larger
The ionic radii of anions are _____________ than their neutral atom The ionic radii of cations are _____________ than their neutral atom Metals closer to the metalloid line have ____________ ionic radii than metals further away. Nonmetals closer to the metalloid line have ___________ ionic radii than nonmetals further away.
s p s d s f
The key feature of representative elements is that they have their valence electrons in the orbital of either _____ or ______ subshells. The key feature of nonrepresentative elements is that they have their valence electrons in the orbital of either ________ and ________ subshells (for transition metals) or _______ and _______ subshells (for lanthanide and actinide elements).
Mercury
The one metal that is not solid at room temperature is ______________, which is liquid at room temperature.
Answer: C. Electronegativity describes how strong an attraction an element will have for electrons in a bond. A nucleus with a larger effective nuclear charge will have a higher electronegativity; Z(eff) increases toward the right side of a period. A stronger nuclear pull will also lead to increased first ionization energy, as the forces make it more difficult to remove an electron. The vertical arrow can be explained by the size of the atoms. As sized decreases, the positive charge becomes more effective at attracting electrons in a chemical bond (higher electronegativity), and the energy required to remove an electron (ionization energy) increases.
The properties of atoms can be predicted, to some extent, by their location within the periodic table. Which property or properties increase in the direction of the arrows shown? I. Electronegativity II. Atomic radius III. First ionization energy A. I only B. I and II only C. I and III only D. II and III only
Positive Negative Positive
The value for ionization energy is always ___________, the value for electron affinity is always _____________. (____________ electron affinities are considered to be inert).
Groups Periods
The vertical columns in the periodic table are called __________ The horizontal rows in the periodic table are called __________
decreases greater decrease
Through the removal of the first electron, electron shielding ______________, causing the valence electron to feel a ___________ attractive force to the nucleus than before. This ______________ in electron shielding causes a massive increase in ionization energy to remove a second electron.
covalent-nonpolar bond covalent-polar bond ionic bond
Two atoms of very similar or same electronegativity that there is very little or no difference in electronegativity are able to form a ________________________ bond. Two atoms that have somewhat differing electronegativities that there is a significant difference in electronegativity are able to form a ______________________ bond. Two atoms that have vastly differing electronegativities that there is a large difference in electronegativity are able to form a __________________________ bond.
Nonmetals
What are the brown colored elements in the periodic table classified as?
-Nuclear Charge -Electron Shielding -Distance between the nucleus and valence electron
What are the factors that affect the difference between first and second ionization energy?
Metals
What are the gray colored elements in the periodic table classified as?
Nonrepresentative elements
What are the green and blue colored elements in the periodic table known as?
Metalloids
What are the green colored elements in the periodic table classified as?
Alkali metals
What are the highlighted elements in the periodic table classified as?
Alkaline earth metals
What are the highlighted elements in the periodic table classified as?
Halogens
What are the highlighted elements in the periodic table classified as?
Noble gases
What are the highlighted elements in the periodic table classified as?
Transition metals
What are the highlighted elements in the periodic table classified as?
Chalcogens
What are the purple-highlighted elements in the periodic table classified as?
Representative elements
What are the yellow colored elements in the periodic table known as?
Answer: C. Atomic radius is determined by multiple factors. Of the choices given, the number of valence electrons does have an impact on the atomic radius. As one moves across a period (row), protons and valence electrons are added, and the electrons are more strongly attracted to the central protons. This attraction tightens the atom, shrinking the atomic radius. The number of electron shells is also significant, as demonstrated by the trend when moving down a group (column). As the electron shells are added that separate the positively charged nucleus from the outermost electrons, the electrostatic forces are weakened, and the atomic radius increases. The number of neutrons is irrelevant because it does not impact these attractive forces.
What determines the length of an element's atomic radius? I. The number of valence electrons II. The number of electron shells III. The number of neutrons in the nucleus A. I only B. II only C. I and II only D. I, II, and III
Germanium's atomic number is 32, thus has a nuclear charge of 32. There are 2 inner electrons on the n=1 valence, 8 inner electrons on the n=2 valence, and 18 inner electrons on the n=3 valence, thus it's electron shielding is 2 + 8 + 18 = 28 Z(eff) = 32 - 28 = +4 Z(eff) = +4
What is the effective charge of germanium?
Lithium's atomic number is 3, thus has a nuclear charge of 3 There are 2 inner electrons shielding the valence electrons Z(eff) = 3 - 2 = +1 Z(eff) = +1
What is the effective charge of lithium?
Phosphorous's atomic number is 15, thus has a nuclear charge of 15. There are 2 inner electrons on the n=1 valence and 8 inner electrons on the n=2 valence, thus it's electron shielding is 8 + 2 = 10. Z(eff) = 15 - 10 = +5 Z(eff) = +5
What is the effective charge of phosphorous?
[Ne] 3s(2) 3p(8)
What is the electron configuration of a chloride (Cl-) anion?
[He] 2s(2) 2p(8)
What is the electron configuration of a sodium (Na+) cation?
[Ar] 4s(2) 3d(4)
What is the electron configuration of an iron (II) (Fe(2+)) cation?
[Ar] 4s(2) 3d(10) 4p(1)
What is the electron configuration of gallium?
[He] 2s(2) 2p(4)
What is the electron configuration of oxygen?
[Ne] 3s(1)
What is the electron configuration of sodium
Answer: C. When n=3, l=0, 1, or 2. The highest value for l in this case is 2, which corresponds to the d subshell. Although the 3d block appears to be part of the fourth period, it still has the principal quantum number n=3. In general, the subshells within an energy shell increases in energy as follows: s < p < d < f (although there is no 3f subshell).
What is the highest-energy orbital of elements with valence electrons in the n = 3 shell? A. s-orbital B. p-orbital C. d-orbital D. f-orbital
Polar covalent
What type of bond is shown between the oxygen and hydrogen atom?
Ionic
What type of bond is shown between the sodium and chloride atoms?
Nonpolar covalent
What type of bond is shown between the two carbon atoms?
Answer: B. Iron is a transition metal. Transition metals can often form more than one ion. Iron, for example, can be Fe(2+) or Fe(3+). The transition metals, in these various oxidation states can often form hydration complexes with water. Part of the significance of these complexes is that, when a transition metal can form a complex, its solubility within the related solvent will increase. The other ions given might dissolve readily in water, but because none of them are transition metals, they will not likely form complexes.
When dissolved in water, which of the following ions is most likely to form a complex ion with H2O? A. Na+ B. Fe(2+) C. Cl- D. S(2-)
Down Left
When going ___________ in the periodic table, you are adding electrons to a higher energy level, which is further away from the nucleus, causing an increase in the atomic radii. When going __________ in the periodic table, the nuclear charge decreases, leading to less attraction of valence electrons to the nucleus, resulting in the valence shell to expand, resulting in an increase in atomic radii.
Ag - B Al - A K - A P - A Pb - A Li - A Pu - B U - B Cu - B Cf - B Zn - B B - A N - A Np - B O - A He - A
Which of the following are representative elements (A) and which are nonrepresentative (B)? -Ag -Al -K -P -Pb -Li -Pu -U -Cu -Cf -Zn -B -N -Np -O -He
Answer: D. The effective nuclear charge refers to the strength with which the protons in the nucleus can pull on electrons. This phenomenon helps to explain electron affinity, electronegativity, and ionization energy. In (A.), the nonionized chlorine atom, the nuclear charge is balanced by the surrounding electrons: 17 p+/17 e-. The chloride ion, (B.), has a lower effective nuclear charge because there are more electrons than protons: 17 p+/18 e-. Next, elemental potassium (C.), has the lowest effective nuclear charge because it contains additional inner shells that shield its valence electron from the nucleus. (D.), ionic potassium, has a higher effective nuclear charge than any of the other options do because it has the same electron configuration as Cl- (and the same amount of shielding from inner shell electrons as neutral Cl) but contains two extra protons in its nucleus; 19 p+/18 e-.
Which of the following atoms or ions has the largest effective nuclear charge? A. Cl B. Cl- C. K D. K+
Answer: B. This question requires knowledge of the trends of electronegativity within the periodic table. Electronegativity increases as one moves from the left to right for the same reasons that effective nuclear charge increases. Electronegativity decreases as one moves down the periodic table because there are more electron shells separating the nucleus from the outermost electrons. In this question, chlorine is the furthest toward the top-right corner of the periodic table.
Which of the following elements has the highest electronegativity? A. Mg B. Cl C. Zn D. I
Answer: B. This block represents the alkaline earth metals, which form divalent cations, or ions with a +2 charge. All of the elements in Group IIA have two electrons in their outermost s subshell. Because loss of these two electrons would leave a full octet as the outermost shell, becoming a divalent cation is a stable configuration for all of the alkaline earth metals. Although some of these elements might be great conductors, they are not as effective as the alkali metals, eliminating (A.). (C.) is also incorrect because, although forming a divalent cation is a stable configuration for the alkaline earth metals, the second ionization energy is still always higher than the first. Finally, (D.) is incorrect because atomic radii increase when moving down a group of elements because the number of electron shells increases.
Which of the following is an important property of the group of elements shaded in the periodic table above? A. These elements are the best electrical conductors in the periodic table B. These elements form divalent cations C. The second ionization energy for these elements is lower than the first ionization energy D. The atomic radii of these elements decrease as one moves down the column
Answer: D. Ionic bonds are formed through unequal sharing of electrons. These bonds typically occur because the electron affinities of the two bonded atoms differ greatly. For example, halogens have high electron affinities because adding a single electron to their valence shells would create full valence shells. In contrast, the alkaline earth metals have very low electron affinities and are more likely to be electron donors because the loss of two electrons would leave them in full valence shells. (A.) states the opposite and is incorrect because the halogens have high electron affinity and the alkaline earth metals have low electron affinity. (B.) is incorrect because equal sharing of electrons is a classic description of covalent bonding, not ionic. (C.) is a true statement, but is not relevant to why ionic bonds form.
Why do halogens often form ionic bonds with alkaline earth metals? A. The alkaline earth metals have much higher electron affinities than the halogens. B. By sharing electrons equally, the alkaline earth metals and halogens both form full octets. C. Within the same row, the halogens have smaller atomic radii than the alkaline earth metals. D. The halogens have much higher electron affinities than the alkaline earth metals.
Distance Two
You determine the atomic radii by finding the _________________ between two identical nuclei and divide by _______.
First ionization energy Second ionization energy
_________ ionization energy is the energy needed to remove one electron from an atom. _________ ionization energy is the energy needed to remove a second electron from an atom after already removing one.
Second First Second First
__________ ionization energy is greater than ___________ ionization energy since it requires more energy to remove the _____________ than the ___________ electron.
Nuclear charge
_____________ ____________ itself does not affect the difference in ionization energies, but it still plays a role in determining it through the calculation of effective nuclear charge.
Effective nuclear charge Nuclear charge Electron shielding Z(eff) = Z - S
_______________ ___________ ___________ is the actual amount of charge from the nucleus that is attracting the valence electrons inward. It is determined by the difference between the _____________ ___________ and ______________ _______________. __________ = ________ - _________
Electron affinity
_______________ _______________ is the energy dissipated by a gaseous species when it gains an electron. X + e- -----> X- + energy
Electron Shielding
________________ ______________ is the phenomenon where electrons from the inner shells of an atom repel the valence electrons from being attracted by the nucleus.
Electronegativity
___________________ is the strength of an atom in a molecule to attract electrons to itself.
Electronegativity Atomic radii Electron affinity Ionization energy
_____________________ increases the further up and the further right you go on the periodic table. _____________________ increases the further down and further left you go on the periodic table. ________________________ (generally) increases the further right and the further up you go on the periodic table. __________________________ (generally) increases the further up and the further right you go on the periodic table.
Electronegativity Electronegativity
______________________ is the tendency for an atom to attract a bonding pair of electrons. ____________________ unlike other periodic trends is measured within a bond.
