Test 5- Periodicity Flashcards Preview

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Flashcards in Test 5- Periodicity Deck (37):
1

Johann Wolfgang Dobereiner

- German chemist who is best known for work that foreshadowed the periodic for the chemical element
- Discovered trends in certain properties of select groups of elements
- Proposed the Law of Triads

2

John Alexander Reina Newlands

- The first person to devise a periodic table of elements arranged in order of their relative atomic weights
- Published in 1865 his law of octaves, which stated that any given element will exhibit analogous behavior to the eighth element following it in the table.

3

Lothar Meyer

- German scientist best known for his part in the periodic classification of the elements
- Organized 28 elements by atomic weight and valence electrons
- Came before Mendeleev- possibly sent to him

4

Dmitri Mendeleev

- Came to him in a dream
- Russian chemist proposed an organization scheme of all the know elements based on valence electrons and atomic weight
- Predicted the existence of 8 new elements and their masses

5

Henry Moseley

In 1911 he determined that each element had a unique atomic number using x-rays

6

Periodic Law

The Periodic Law states that the physical and chemical properties of the elements recur in a systematic and predictable way when the elements are arranged in order of increasing atomic number.

7

Characteristics of Metals

- Metals are malleable, ductile, can conduct heat and electricity, have a metallic luster, have a few valence electrons (1-3), exhibit metallic luster, can be representative, transitional, or inner transitional (rare earth)
- Valence electrons not attracted to nucleus instead of attaching to 1 atom, floats between all atoms, bonding them together, 1 VE not as strong (metallic bond), number of VE determines strength of metallic bond

8

Delocalized

float among all electrons and cause metallic bond

9

Malleable

can hammer into shape

10

Ductile

pulled in wires

11

Representative/Transitional/Intertransitional

- Representative metals- Always lose the same amount of electrons when they lose electrons, form ions that always have the same charge
- Transitional and inter transitional form ions with varying charges

12

Characteristics of nonmetals

- Opposite of metals
- not malleable
- not ductile
- doesn’t conduct heat and electricity
- many Valence electrons

13

Can determine the __ from the periodic table

electron configuration

14

Liquids on the periodic table

Bromine (not liquid) and Mercury

15

Gases on the periodic table

- 7 diatomic elements (Br, I, Cl, N, H, O, F) Br and I aren’t gases but they’re atomic
- Nobel gases.

16

Solids on the periodic table

Majority of elements, mostly metals. Metalic doesn’t equal Solid

17

Diatomic Elements

have to find another element to hook up with or even itself. Have to write little 2 after each one unless in compound because there’s at least 2 of them

18

Families or Groups

- Indicates number of valence electrons
- Valence electrons are responsible for chemical properties or elements. Ex. Outside ring or Bohr model

19

Alkali Metals

- Column 1 or 1A (except H)
- Shiny, soft, low density (will float), never found in elemental state in nature, one valence electron
- Loses one VE to form an ion with a +1 charge
- Reacts easily with oxygen
- Produces hydrogen and a solution in exothermic reaction with water (produces heat, basic substance left behind- bases can be just as harful as acids)
- Need the ion in your diet

20

Alkaline Earth Metals

- Similar to alkali metals- just not as
- 2 valence electrons, loses these two valence electrons to form an ion with a +2 charge
- Column 2 or 2A
- Not as vigorously react
- Not as shiny, soft, etc
- A little harder, more dense, etc

21

Halogens

- All nonmetals
- Column 7 or 17
- 7 Valence electrons
- Gains one electron to form ions with a -1 charge
- Diatomic

22

Nobel or Inert gases

- Chemically inactive, rarely form compounds
- Column 8 or 18
- Has 8 Valence electrons (except He which has only 2 electrons total)
- Stable

23

Period or Series

- The family or group indicates the number of VE, the period or series indicates the number of total energy levels of electrons an atom possesses
- Row number=number of energy levels
- Find number of electrons by counting how many over from the beginning of sublevel section, ex. Ba is 2 over
- In transitional metal section, total number of energy levels is different from what electron level the last electron is found
- D section- EL= n-1
- F section- EL=N-2
- Three series- Lanthanide, Actinide, and Glen Seaborg (intertransitional elements at the bottom)

24

Sublevels

- S includes the Alkali and Alkaline earth metals
- P block includes the nonmetals
- D block includes the transitional metals
- F block includes the inner transitional or rare earth metals
- Number of rows it takes up=number or electrons it can hold

25

Atomic Radius

- half the distance between the nuclei of identical atoms that are bonded together
- Determines all properties

26

Smaller radius

greater attractive force and easier to gain electrons

27

Atomic Radius trends

- increase down a group because of increasing energy levels
- decreases across a period due to the increasing positive charge of the radius

28

Shielding effect

The inner electrons shied the outer electrons from the attractive force of the nucleus. Inner electrons=repelling force

29

Columbic Attraction

The attraction between oppositely charged particles

30

Further down row=

greater radius and shielding effect- larger electron level and more inner electrons

31

How elements react with water depends on

- How tightly the atom holds its Valence electrons.
- Ex. Li and K both have 1 VE to react but Li electron is closer and harder to remove- K electron is farther and easier to remove and very reactive

32

Cations vs. Anions

- Cation is always has a smaller atomic radius than parent atom since it loses and electron and the columbic attraction increases
- Anion always larger atomic radius than parent atom- columbic attraction decreases

33

Ionization energy

- The energy required to remove AN electron from a neutral atom of an element in the gaseous state
- A+energy (removing electron) → A+ (+ because electron taken away) + e- (electron released)
- Measured in KJ
- The ionization energy increases as the atomic radius decreases
- Larger radius=easier to remove electrons
- Takes electrons from the highest EL

34

Metallic Characteristics

- The easier an element loses electrons, the more metallic it is considered
- Therefore, the further left and down on the periodic table, the more metallic in character
- Metals lose electrons easily
- Fr is the most metallic and has the largest radius
- Nonmetals gain electrons easiest- smallest radius
- Not He because it’s a nobel gas and doesn’t react so it F
- Don’t count nobel gases when talking about metallic/nonmetallic
- Fluorine loves electrons
- Fr doesn’t want electrons at all

35

Electron Affinity

- The energy change that occurs when an electron is acquired by a neutral atom
- Most atoms release energy when they acquire electrons
- A+ + e- → A + energy
- Some atoms must be forced to gain electrons (Fr)
- A+ + e- + energy → A
- The more nonmetallic the higher the electron affinity
- Measured in KJ

36

Ionic Radii

- Positive ions- Lose electrons- Cations
- Cations are smaller than the parent atom
- Negative Ions- Gain electrons- Anions
- Anions are larger than the parent atom
- Periodic trend- Ionic radius decreases from left to right as nuclear charge increases on cation until anions are reached and then the radius sharply increases due to addition of energy levels with anion formation
- Group trend- Ionic radius increases down a groups as number of energy level increase

37

Electron Negativity

- Linus Pauling
- Relative scale of 0-4
- Measure of the ability of an atom in a compound to attract electrons from another atom in a compound
- More nonmetallic, the electronegativity increase