Electrons in Atoms A chapter dedicated just for you electrons

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Presentation transcript:

Electrons in Atoms A chapter dedicated just for you electrons Electrons in Atoms A chapter dedicated just for you electrons! Chapter 5

Light Waves Amplitude = wave’s height Wavelength = wave’s distance Frequency = wave cycles (hertz)

Electromagnetic radiation = A big word for Light

Electromagnetic Radiation Waves have a frequency Use the Greek letter “nu”, , for frequency, and units are “cycles per sec” or Hertz (Hz) All radiation:  •  = c where c = velocity of light = 2.998 x 108 m/sec

Electromagnetic Spectrum Long wavelength --> small frequency INVERSE Short wavelength --> high frequency RELATIONSHIP

Electromagnetic Spectrum In increasing energy, ROY G BIV

White light is actually all the colors together A prism bends the different colors to where we can see them

Atomic Emission Spectra and Niels Bohr Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the ATOMIC EMISSION SPECTRA of excited atoms - Remember the Bohr-rings! Niels Bohr (1885-1962)

Atomic Emission Spectra Excited Atoms Excited atoms emit light of only certain wavelengths The wavelengths of emitted light depend on the element Each element has its own Atomic Emission Spectra… like a FINGERPRINT ELECTRICITY or HEAT EXCITED ELECTRONS!

5.2 = More about electrons Heisenberg Uncertainty Principle Problem of defining nature of electrons in atoms solved by W. Heisenberg Cannot simultaneously define the position of an electron We define e- energy exactly but do not know exact position W. Heisenberg 1901-1976

We know area the propeller is, but we cannot locate its exact position

Arrangement of Electrons in Atoms Electrons in atoms are arranged as LEVELS ORBITALS SUBLEVELS

QUANTUM NUMBERS n (principal) ---> energy level = FROM BOHR! The shape, size, spin, and energy of each orbital is a function of 4 quantum numbers which describe the approximate location of an electron n (principal) ---> energy level = FROM BOHR! l (orbital) ---> shape of orbital ml (magnetic) ---> designates a particular suborbital s (spin) ---> spin of the electron (clockwise or counterclockwise) Think of the 4 quantum numbers as the address of an electron… Country > State > City > Street

QUANTUM NUMBERS So… if two electrons are in the same place at the same time, they must be repelling (like charges repel!) The Pauli Exclusion Principle says that an atomic orbital may describe at most 2 electrons. If two electrons are in the same energy level, the same sublevel, and the same orbital, they must repel = spin (S) Pauli = spinning electrons (repel)

Energy Levels Each energy level has a number called the PRINCIPAL QUANTUM NUMBER, n (energy level) Currently n can be 1 thru 7, because there are 7 periods on the periodic table The period #s = energy levels

Energy Levels n = 1 n = 2 n = 3 n = 4

Types of Orbitals The most probable area to find these electrons takes on a shape So far, we have 4 shapes They are named s, p, d, and f No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise

Types of Orbitals (l) s orbital p orbital d orbital

Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

Aufbau’s Principal Electrons occupy the orbitals of lowest energy first! Aufbau = electron order…. Least to most energy!

Orbitals and the Periodic Table Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) s orbitals d orbitals p orbitals f orbitals

Electron Configurations A list of all the electrons in an atom Must go in order (Aufbau principle) 2 electrons per orbital, maximum (Pauli Ex.) We need electron configurations so that we can determine the number of electrons in the outermost energy level These are called valence electrons The number of valence electrons determines how many and what this atom can bond to in order to make a molecule 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

Electron Configurations 2p4 Number of electrons in the sublevel Energy Level (n) Sublevel 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

Diagonal Rule 1 2 3 4 5 6 7 s s 2p s 3p 3d s 4p 4d 4f s 5p 5d 5f 5g? Steps: Write the energy levels top to bottom. Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level. Draw diagonal lines from the top right to the bottom left. To get the correct order, follow the arrows! 1 2 3 4 5 6 7 s s 2p s 3p 3d s 4p 4d 4f By this point, we are past the current periodic table so we can stop. s 5p 5d 5f 5g? s 6p 6d 6f 6g? 6h? s 7p 7d 7f 7g? 7h? 7i?

Let’s Try It! Write the electron configuration for the following elements: H 1s1 Li 1s2 2s1 N 1s2 2s2 2p3 Ne 1s2 2s2 2p6 K 1s2 2s2 2p6 3s2 3p6 4s1 Zn 1s2 2s2 2p6 3s2 3p6 4s2 3d10

Orbital Diagrams Graphical representation of an electron configuration One arrow represents one electron Shows spin and which orbital within a sublevel

Orbital Diagrams Hund’s Rule In orbitals of EQUAL ENERGY (p, d, and f), place one electron in each orbital before making any pairs All single electrons must spin the same way I nickname this rule the “Monopoly Rule” In Monopoly, you have to build houses EVENLY. You can not put 2 houses on a property until all the properties have at least 1 house.

Lithium Group 1A Atomic number = 3 1s22s1 ---> 3 total electrons

Carbon Group 4A Atomic number = 6 1s2 2s2 2p2 ---> 6 total electrons HUND’S RULE! Same energy level = 1 electron until All filled

Draw these orbital diagrams! Lithium (Li) Beryllium (Be) it helps to do the Boron (B) e- configuration Carbon (C) first Nitrogen (N) Neon (Ne) Sodium (Na)

Ch. 6.1 & 6.2 Dmitri Mendeleev (1869) Mendeleev published classification schemes for elements known to date. The periodic table has organized the similar properties and reactivities by certain elements Later, Henri Moseley established that each elements has a unique atomic number

The Periodic Table http://www.chemsoc.org/viselements/pages/periodic_table.html

Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties

Periodic Table Expanded View The Periodic Table can be arrange by subshells The s-block is Group IA and & IIA The p-block is Group IIIA - VIIIA The d-block is the transition metals The f-block are the Lanthanides and Actinide metals

Periodic Table: Metallic arrangement Layout of the Periodic Table: Metals vs. nonmetals

METALS 80 percent of all elements Are GOOD conductors of heat and electric current (Give up electrons) Have a high luster or sheen Reflects light All metals are solids at room temperature (except mercury Hg) They are ductile (can be drawn into wires) They are malleable (can be hammered down)

NONMETALS Upper-right corner Variation in physical properties: Accept electrons Most are gas at room temperature (N and O) Few are solids (sulfur and phosphorus) One is a liquid (bromine—dark red) Most nonmetals have properties opposite of metals: Poor conductors (except carbon) Brittle when solid

METALLIODS Hug the stair-step line Have properties similar to both metals and nonmetals, depending on certain conditions Give up or accept electrons EXAMPLE: Pure silicon is a poor conductor of electrical current But… if silicon is mixed with boron, it is a great conductor of electrical current Used to make computer chips

Across the Periodic Table Periods: Are arranged horizontally across the periodic table These elements have the same number of valence shells. 2nd Period 6th Period

Down the Periodic Table Family: Are arranged vertically down the periodic table = These elements have the same number electrons in the the valence shell

Groups The specific colors of the squares distinguish the groups of elements Group 1A = Alkali Metals Group 2A = Alkaline Earth Metals Alkali is the arabic word al aqali which means wood ashes Wood ashes are rich in Na and K

Groups The specific colors of the squares distinguish the groups of elements Group 7A = Halogens Halogen comes from the Greek word hals (salt) and the Latin word genesis (to be born) Chlorine, bromine, and iodine can be prepared from their salts

Electron Configurations Elements can be sorted into noble gases, representative elements, transition metals, or inner transition metals based on their electron configurations The noble gases are group 8A… they rarely take place in reactions IF YOU LOOK AT THEIR CONFIGURATIONS… THERE VALENCE SHELL IS FULL!!!

Infamous Families of the Periodic Table Notable families of the Periodic Table

Representative Elements Groups 1A- 7A They display a wide range of physical and chemical properties For representative elements… The group number is equal to the number of electrons in the highest occupied energy level

Transition Elements Transition metals are characterized by electrons in the d sublevel Inner transition metals are characterized by electrons in the f sublevel The periodic table is organized into ‘blocks’ Based on sublevels (except for helium)

General Periodic Trends 6.3 Atomic and ionic size Ionization energy (electronegativity– learn in future chapter)

Atomic Radius Is one-half the distance between the nuclei of two atoms of the same element when atoms are joined

Atomic Size Size increases- going down a group Each element has 1 more proton and 1 more electron than the preceding one The increasing nuclear charge pulls the electrons in the highest occupied energy level closer to the nucleus and the atomic size decreases Size decreases- going across a period

Which is Bigger? Na or K ? K Na or Mg ? Na Al or I ? Al

What are IONS?? An atom or group of atoms that has a positive or negative charge An atom is electrically neutral because it has equal numbers of protons and electrons… EX: Na has 11 protons and 11 electrons = 0 charge Positive (cation) and negative (anion) ions form when electrons are transferred between atoms ELECTRONS determine a charge!

IONS Cations = positive charge Lose electrons Usually in metallic elements Anions = negative charge Gain electrons Usually nonmetallic elements How do you think the trend is going go?

Ion Sizes Does the size go up or down when losing an electron to form a cation?

Ion Sizes CATIONS are SMALLER than the atoms from which they come Li + , 78 pm 2e and 3 p Forming a cation Li,152 pm 3e and 3p CATIONS are SMALLER than the atoms from which they come The electron/proton attraction has gone UP and so size DECREASES (forced closer to nucleus)

Ion Sizes Does the size go up or down when gaining an electron to form an anion?

Ion Sizes Forming an anion - , 133 pm 10 e and 9 p F, 71 pm 9e and 9p Forming an anion ANIONS are LARGER than the atoms from which they come The electron/proton attraction has gone DOWN and so size INCREASES

Cl or Cl2+ ? Cl K+ or K ? K F or F-? F- I- or Br- ? I- (b/c size) Which is Bigger? Cl or Cl2+ ? Cl K+ or K ? K F or F-? F- I- or Br- ? I- (b/c size)

Ionization Energy IE = energy required to remove an electron from an atom (in the gas phase) This is called the FIRST ionization energy because we removed only the OUTERMOST electron

Trends in Ionization Energy IE increases across a period because the positive charge increases. Metals lose electrons more easily than nonmetals Nonmetals lose electrons with difficulty (they like to GAIN electrons)

Trends in Ionization Energy IE increases UP a group Because size increases

Which has a higher 1st ionization energy? Mg or Ca ? Mg Al or S ? S Cs or Ba ? Ba

Electronegativity Is a measure of the ability of an atom in a molecule to attract electrons Will talk about more later…