Periodic Law
Halogens properties
Most reactive nonmetal React vigorously with most metals to form salts Gain 1 electron to be stable (opposite A1) They react vigorously with most metals to form salts.
Noble gases properties
Most unreactive nonmetal Always found in nature in element form except Xe
Based on what you know would Na or Ne have more nuclear charge/ positive charge/ attraction/ pull
Na (Ne?) because more electron and protons to pull
Alkaline-earth metals.
The elements of Group 2 of the periodic table (beryllium, magnesium, calcium, strontium, barium, and radium)
atom's electron affinity.
The energy change that occurs when an electron is acquired by a neutral atom
ionization energy, IE (or first ionization energy, IE1)
The energy required to remove one electron from a neutral atom of an element
Atomic radius
one-half the distance between the nuclei of identical atoms that are bonded together
periods
organized horizontally in rows, (There are a total of seven periods of elements in the modern periodic table.)
cation
positive ion
Group 3-12=
Inner transition metals
halogens
The elements of Group 17 (fluorine, chlorine, bromine, iodine, and astatine)
anion
negative ion
What type of ions do metals form
+
What kind of ions do group 1 usually form? Why?
+ Lose 1 electron same proton: more positive charge
What is the relationship of positive ions the parent and negative ions
1+< parent< 1-
Why is group 2 still reactive
2 valence electron from being stable
How many periods and groups
7 periods and 18 groups
Why are noble gas unreactive
8 valence electron
Consider 2 main group elements A and B. A has an ionization energy1 of 419 kJ/mol. B has an ionization energy1 of 1000kJ/mol. A) in which block can you find element A and B B) which element is more likely to form a positive ion
A s block B p block A
The noble gases have A)high organization energy B) high electron affinity C) large atomic radii D) A tendency to for both cations and anions
A)high organization energy
Which group is the most reactive? Which is the least?
A1 A7 and A8
Group 1=
Alkali metals
Group 2=
Alkaline earth metals
ionization.
Any process that results in the formation of an ion
Why was Mendeleev able to predict future elements? Where they accurate
Arranged by atomic mass and saw a pattern Yes
What is the most rare element in the world
At
Which two elements are more likely to have the same charge on there ions A) Se and As B) Sn and Si C) Ca and Rb D) I and Xe
B) Sn and Si
Why is K less strong than Na? What can you infere
Because inner electron cancel out postivie nuclear traction as you go down Farther distance Because K is weaker and it can lose electron easy because not as strongly attracted to nucleus
Why is hydrogen not a alkali metal? Then why is it in group 1 at all?
Because it isn't a metal Because it has 1 valence electron
Which blocks have a consistent pattern
Block s and block p
Of the element Ga, Br, and Ca which has the highest electronegativity? Explain the trend
Br because as you go to the left and up you have higher electronegativity. It has a strong nuclear charge that isn't as canceled out as other elements below it
Which of the following elements has the highest eletronegativity A) O B) H C) F D) C
C) F
"Special" quality about group 14
Can lose/gain 4 electron (Not usually an ion but turn into a different compound)
Trends for ionic radii (2)
Cation Period trend:Cationic radii decrease across a period, because the electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level. Cation Group trend:The outer electrons in both cations and anions are in higher energy levels as one reads down a group. Thus, just as there is a gradual increase of atomic radii down a group, there is also a gradual increase of ionic radii. Anion Period trend:Anionic radii decrease across each period for the elements in Groups 15-18. The reasons for this trend are the same as the reasons that cationic radii decrease from left to right across a period. Anion Group trend:The outer electrons in both cations and anions are in higher energy levels as one reads down a group. Thus, just as there is a gradual increase of atomic radii down a group, there is also a gradual increase of ionic radii.
Explain trends for iodic radii in your own words
Cation/anion period: The electron cloud become smaller because the remaining electrons are drawn closer to the nucleus by its unbalance positive charge Cation/anion group: gradual increase of ionic radii down a group because electrons are in higher energy levels as you go down to group
Why is ionization energy2 higher than ionization energy1, ionization energy3 higher than ionization energy2, ...
Due to the fewer electron remaining within the atom to shield the attractive force of the nucleus Protons are the same but less electron to cancel out nuclear attraction
What is the difference throughout group 1
Inner electron increase as you go down
Groups
Elements are arranged vertically in the periodic table
What is Moseley known for? How is that related to today?
Elements fit into patterns better when they were arranged according to atomic number That is what we use today, this modern periodic table
Strongest and weakest electronegativity
Fluorine Francium
What halogen is the most reactive in terms of electron affinity? Which is the least reactive in terms of electron affinity?
Fluorine Francium
What happens when group 1 loses an electron? Explain
Forms an ion losing 1 electron but having the same proton No longer nuetral because there is a net change in charge
The more negative electron affinity that means
Greater attraction for an electron ❤️
Which group would have the highest electron affinity? Why
Group 17/ halogens because they want to get 1 electron
What is stronger H or Fr
H
Give an example of electronegativity
H20 O is stronger and pulls H close even if they share electron evenly
Group 17=
Halogens
Properties of p block
Harder than the s block alkaline earth metals but softer and less dense than the d block metals Soft and brittle vary greatly because metals, nonmetals, metaloids
Give properties of alkali metals
In there pure state, all of the alkali metals have a silvery appearance and are soft enough to cut with a knife Solid Too reactive to be found in nature in element form HYDROGEN IS NOT INCLUDED They combine vigorously with most nonmetals. And they react strongly with water to produce hydrogen gas and aqueous solutions of substances known as alkalis. alkali metals are usually stored in kerosene. Proceeding down the column, the elements of Group 1 melt at successively lower temperatures.
What were noble gas first called? Why change?
Inert (unreactive) gases Wait Xe is reactive Noble gas
Ion+ion
Ionic compound
How can you measure the radius of an atom
It's hard to because it's fuzzy and can vary depending on the conditions Measure half of the distance between the nuclei of identical atoms that are bonded together
Knowing the group means
Knowing the characterlistics
F block=
Lanthanides Actinides
Lanthanides and actinides properties
Lanthanides- shiny metals similar in reactivity to the group 2 alkaline earth metals Actinides- radioactive
First scientist we talked about to organized periodic table
Mendeleev
periodic
Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals. repeating pattern
Inner transition metals properties
Metals with typical metallic properties Often referred as transition metals Not as reactive as A1-A2 Still reactive good conductors of electricity have a high luster. less reactive than the alkali metals and the alkaline-earth metals. Some are so nonreactive
Would metals increase in radii size or decrease in radii size? Nonmetals? Why
Metals: decrease because they want to lose electrons Nonmetal: increase because they want to gin electrons
Carbon is a __
Molecular compound
The jump is usually how much bigger
More than 10x
Of the elements Mg, Cl, Na, and P which has the largest atomic radii? Explain the trend
Na has least amount of positive charge from any increasing protons. In comparison the the rest it is the least attracted to the nucleus (trend across the period)
Group 18=
Noble gases
P block=
Nonmetals except H and He Metalliods
What does the length of each period determine
Number of electrons that can occupy the sublevels being filled in that period
What is Mendeleev known for? How did He do it? Explain
Organized the first periodic table Arranged the elements in order of increasing atomic mass Left open space for elements He didn't know He predicted the exsistance, properties, and characerlistics of empty spaces
Which block needs more ionization energy? Why
P block because as you go across there is stronger nuclear attraction
Explain the trends for electron affinity in your own words
Period and group: increase both across to the right add up the group because there is an increasing in protons and therefore nuclear charge. The strong nuclear charge will attract other electrons more strongly.
Trends for ionization energy
Period trend:In general, ionization energies of the main-group elements increase across each period. This increase is caused by increasing nuclear charge. A higher charge more strongly attracts electrons in the same energy level. Increasing nuclear charge is responsible for both increasing ionization energy and decreasing radii across the periods Group trend:Among the main-group elements, ionization energies generally decrease down the groups. Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the nucleus. Therefore, they are removed more easily. Also, as atomic number increases going down a group, more electrons lie between the nucleus and the electrons in the highest occupied energy levels. This partially shields the outer electrons from the effect of the nuclear charge. Together, these influences overcome the attraction of the electrons to the increasing nuclear charge.
Trends for atomic radii
Period trend:The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus (As electrons add to s and p sublevels in the same main energy level, they are gradually pulled closer to the more highly charged nucleus. This increased pull results in a decrease in atomic radii. The attraction of the nucleus is somewhat offset by repulsion among the increased number of electrons in the same outer energy level. As a result, the difference in radii between neighboring atoms in each period grows smaller) Group trend:In general, the atomic radii of the main-group elements increase down a group (As electrons occupy sublevels in successively higher main energy levels farther from the nucleus, the sizes of the atoms increase)
Trends for electron affinity
Period trend:as electrons add to the same p sublevel of atoms with increasing nuclear charge, electron affinities become more negative across each period within the p-block. Group trend:electrons are added with greater difficulty down a group. This pattern is a result of two competing factors. The first is a slight increase in effective nuclear charge down a group, which increases electron affinities. The second is an increase in atomic radius down a group, which decreases electron affinities. In general, the size effect predominates.
Trends for electronegativity
Period trend:electronegativities tend to increase across each period, although there are exceptions Group trend:Electronegativities tend to either decrease down a group or remain about the same.
Explain trends for ionization energy in your own words
Period: increase across to the right of the periods because increase nuclear charge which attracts electron more strongly therefore needing more energy to remove an electron Group: decrease down the group because as you go down there are high energy levels which are farther from the nucleus and the nuclear charge is canceled out from the inner electron making it easier to remove electron at the bottom of the periodic table
Put the trends for atomic radii in your own words
Period: smaller in size across to the right because more protons and therefore an increase in positive nuclear charge. This increase pulls/attracts electron toward the nucleus Group: increase in size down the period because more electrons to fill more energy sublevels/shells farther from the nucleus. This increase in electrons increase the size of the electron cloud
Explain the trends for electronegativity in your own words
Periods: increase as you go right because more protons and therefore nuclear charge to pull in other electron Group: decrease do the group because more electron to shield the nuclear attraction
Consider 4 hypothetical main group elements Q, R, T, X that have the outer electron configuration: Q: 3s^2 3p^5 R: 3s^1 T: 4d^10 5s^2 5p^5 X:4d^10 5s^2 5p^1 A) identify the block location of each element B) which of these elements are in the same period? Same group? C) which element would you expect to have the highest ionization energy1 and lowest? D) which element would you expect to have the highest Ionization energy2 E) which of the elements is most likely to form a 1+ ion
Q: p R: s T: p X: p Same period: Q and R, T and X Same group: Q and T Highest: Q Lowest: X R R
When gaining an electron what happens to energy? What is it value
Release energy Negative
What does +kJ/mole mean? What values can it be assigned
Require energy this is unnatural and happens when it's forced +# and 0
How is Na+ related to Ne
Same electron configuration Same pulling effect
Why are elements placed in the same group
Share similar chemical properities
What do inner electron act like
Shield
Properties of alkaline earth metals
Similar to group 1 but less reactive Still reactive that you can't find them in nature in element form harder, denser, and stronger than the alkali metals. They also have higher melting points.
Louis dot structure
Simple way to draw valence electron
What stays the same across to the right of a period? What changes
Stays: inner electron Distance from nucleus Power of nuclear power? Change: valence electron changes More proton Increasing nuclear attraction Decrease in size
transition elements
The d-block elements are metals with typical metallic properties
valence electrons
The electrons available to be lost, gained, or shared in the formation of chemical compounds
alkali metals
The elements of Group 1 of the periodic table (lithium, sodium, potassium, rubidium, cesium, and francium)
main-group elements
The p-block elements together with the s-block elements
periodic law
The physical and chemical properties of the elements are periodic functions of their atomic numbers when the elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals.
When atoms are forced to gain an electron what is its condition? What elements have this
These ions are unstable and lose the added electron spontaneously Noble gases and group 2 (stable in terms of electron configuration)
What is the purpose of a compound
To gain stability
What will halogens never gain
Two electrons/2- because it's impossible to gain another electron after assuming a noble gas configuration
What does reactivity depend on
Valence electron
What stays the same in groups as you go down? What changes
Valence electron Inner electron
Why is the ionization energy higher is A2 then A1
Valence electron is closer to the nucleus because of increasing nuclear charge
Why is group 1 so reactive
Wants to be in a stable form Having 8 electrons is stable Wants to have 8 electron/octet/noble gas configuration/full/stable Losing 1 electron would make it stable IT IS DYING TO LOSE ONE ELECTRON IT'S DYING TO REACT
Why is there a huge jump in ionization energy sometimes? Explain
When ion assumes a noble gas configuration (is stable)
periodic table
arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.
ion
atom or group of bonded atoms that has a positive or negative charge
metalloids
boron, silicon, germanium, arsenic, antimony, and tellurium
Electronegativity
measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound.
lanthanides
the 14 elements with atomic numbers from 58 (cerium, Ce) to 71 (lutetium, Lu).
actinides
the 14 elements with atomic numbers from 90 (thorium, Th) to 103 (lawrencium, Lr).
Nonmetals __ metals
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