Periodic Trends And Atomic Theory Flashcards
How is atomic size and energy level reflected on the periodic table?
As you go down a group the energy level increases as the elements have more electron shells. This causes less attraction and a bigger are of the outer shells because of repulsion. Across a period to the right, size decreases because there are more protons which makes more attraction.
Define atomic size
The atomic size of an atom also known as the atomic radius refers to the distance between an atom’s nucleus and it’s valence electrons. Remember, the closer an electron is to the nucleus, the lower it’s energy and the more tightly it is held.
How are ionization and electron affinity alike?
The process of gaining or losing an electron requires energy. There are two common ways to measure this energy change: ionization energy and electron affinity.
Define electron affinity
The change in energy that accompanies the addition of an electron to an atom in the gaseous state. Added into the outer energy shell.
Electron Affinity decreases going down a group because more shells means less nuclear attraction: resulting in a smaller change in energy.
It increases going to the right because the number of valence electrons increases resulting in a bigger “want” of electrons. *group 1+2 have low electron affinity because they are less likely to take electrons electrons and more likely to give them.
Define ionization energy
Anions(non-metals on the right) gain electrons from the cations (metals on the left) that lose their electrons.
The higher the ionization energy, the less likely the atom is to give away an electron, the more likely they are to gain an electron from a low ionization energy leveled atom.
Ionization energy increases going up a group and to the right of the periods.
What is the patterns based on electrons?
Main group pattern: The group number an atom is in is related to the amount of electrons that are in the valence shell.
Define electronegativity
Electronegativity refers to the ability of an atom to attract the electrons of another atom to it when those two atoms are associated through a bond. It is based on an atom’s ionization energy and electron affinity. For that reason electronegativity follows similar trends as it’s two constituent measures.
Use the chart to figure out the electro negativity number, the bigger atom is subtracted by the smaller atom;
-0->0.5( pure covalent)
-0.5->1.7(polar covalent; dipole, unequal sharing of electrons)
-1.7->3.3(mostly ionic)
Compare and contrast ion, isotope and radioisotope
For an ion the electron number varies.
For and isotope the mass will differ because the number of neutrons change.
For an radioisotope they are just isotopes with unstable nuclei.
Ionic vs covalent bonds
Atoms in a covalent bond will share the outer electron. Because of their bond, relatively high energies are required to break them (ex: extremely hot temperatures, and extremely low temperatures). The stronger the electronegativity on one of the atoms the more pull it’ll have on the electron(s) making the bond somewhat polar. This pull goes gets stronger and stronger depending on the atoms, eventually the electro negativity will be so strong for one of the elements it’ll cause the electron to fully break off the one element and be transferred onto the high electronegative atom which is an ionic bond.(ex water with a strong polarity so oxygen, not as high or low temperatures to break them)
Ionic compound
Give and receive electrons.
State at room temperature: crystalline solid
Melting point: high
Electrical conductivity:yes(except when in solid state)
Solubility in water: most are high
Conducts electricity when dissolved in water: yes
Covalent compound
Shares electrons.
State at room temperature: liquid, solid, gas
Melting point: low
Electrical conductivity:no
Solubility in water: most are low
Conducts electricity when dissolved in water: not usually
When and how to draw Lewis structures
Use Lewis structures for ionic bonds.
Write the letter(s) for the atoms down side by side with some space, use dots to represent the outer electrons and draw them around the letters(2 per side). Then draw an arrow from the transferring electron over to it’s place with the new atom. Then beside it write out the letters again surrounded by brackets, only one of the atoms should have electrons surrounding it, beside draw the - and the amount of electrons it gave up and a + with the amount of electrons it gained.
How to draw Lewis structures for covalent bonds
Write the letters that represent the atoms with dots for the valence electrons around the letter. Circle two atoms(one from each atom that will be shared) until all have partners. Beside write out the letters for each atom, draw lines in between the two atoms, one for each shared pair. Then on the opposite side facing away from the other atom draw the electrons that aren’t shared.
Forces- intermolecular vs intramolecular
Intramolecular forces are strong covalent bonds that hold atoms together.
Intermolecular forces are weak relative to covalent bonds, they are the forces between the covalent compounds that hold them all together.
Non-metals tend to form ionic bonds with metals. Non-metals tend to form covalent bonds with other non-metals. How do metals bond to each other?
Metallic bonds are held together by an electrostatic attractive forces between the delocalized electrons called conductive electrons.
