Average Atomic Mass In looking at the masses of the elements on the periodic table it is evident that most values listed aren’t close to being whole numbers Why? These values are actually weighted averages that represent all the naturally occurring isotopes and the relative abundance in which they are found in nature
The proton with a stadardized charge of +1, a mass 1.673 x 10-24 g, and located in the nucleus
Atomic Weight (Average Atomic Mass) is the weighted average mass of all the naturally occurring isotopesof that element.
Average Atomic Mass • The average atomic mass may be determined by calculation from any analysis yielding the actual masses of the isotopes and some relative proportion of their presence (an actual number of nuclides or aspercentages of nuclides) • A “rough estimate” may be obtained if the mass numbers are utilized
classification of matter
example
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pure substances
Pure substances have constant composition. – Elements – Pure substance that CANNOT be broken down into simpler substances by chemical changes. • Consist of one type of element.
Examples: Gold (Au), Phosphorus (P4), Oxygen (O2)
elements
Elements are pure substances that are made up of homoatomic molecules or individual atoms of the same kind
Examples: oxygen gas made up of homoatomic molecules and copper metal made up of individual copper atoms
- The number of atoms in one mole of any element is called Avogadro's number and is equal to 6.022x10^23 . – A one-mole sample of any element will contain the same number of atoms as a one-mole sample of any other element. – One mole of any element is a sample of the element with a mass in grams that is numerically equal to the atomic weight of the element
Compound
Compounds – Pure substances that CAN be broken down into simpler substances by chemical changes. • Consist of two or more types of elements chemically bonded • The properties of compounds are different from the uncombined elements making up the compound. Pure substances and mixtures ^
Examples: H2O, C6H12O6, AgCl
Compounds are pure substances that are made up of heteroatomic molecules or individual atoms (ions) of two or more different kinds
Examples: pure water made up of heteroatomic molecules and table salt made up of sodium atoms (ions) and chlorine atoms (ions)
rows
periods
Rows are called Periods • Seven Seven periods for the seven energy levels(rings)
Formulas are used to represent compounds.
A compound formula consists of the symbols of the elements found in the compound. Each elemental symbol represents one atom of the element. If more than one atom is represented, a subscript following the elemental symbol is used.
if carbon disulfide contains one atom of carbon for every two atoms of sulfur, what is the chemical formula for carbon disulfide?
CS2
• Carbon monoxide, CO – one atom of C – one atom of O • Water, H2 O – two atoms of H – one atom of O • Ammonia, NH 3 – one atom of N – 3 atoms of H
Subtopic
elements
• A substance composed of a single kind of atom. • Cannot be broken down into another substance by chemical or physical means.
Element is the simplest form of matter that can not be broken down by chemical means; One element can be changed into another element by nuclear methods • Are the building blocks of matter • Currently: Type of matter composed of atoms that all have the same atomic #s (Identical atoms) • 118 elements known today Of the 118 known only the first 98 are known to occur naturally on earth • Those that do not occur naturally have been artificially produced by man as synthetic products of nuclear reactions such as Einsteinium , Nobelium
Hydrogen (H) The lightest & the most abundant element in the universe [75%, followed by Helium 23%] • Carbon (C) The 2 nd most abundant element[18.5%] in human body after Oxygen • Oxygen (O) The most abundant element on earth crust [47% followed by Si 28%]
O 65.0 % K 0.34 C 18.5 S 0.26 H 10.0 Na 0.14 N 3.0 Cl 0.14 Ca 1.4 Fe 0.004 P 1.0 Zn 0.003 Mg 0.50 Trace Elements: As, Cr, Co, Cu, F, I, Mn, Mo, Ni, Se, Si, V
A symbol is assigned to each element. The symbol is based on the name of the element and consists of one capital letter or a capital letter followed by a lower case letter. – Some symbols are based on the Latin or German name of the element.
Select the correct name for each: A. N 1) neon 2) nitrogen 3) nickel B. P 1) potassium 2) phogiston 3) phosphorus C. Ag 1) silver 2) agean 3) gold ^
Select the correct name for each: A. N 2) nitrogen B. P 3) phosphorus C. Ag 1) silver^
Select the correct symbol for each: A. Calcium 1) C 2) Ca 3) CA B. Sulfur 1) S 2) Sl 3) Su C. Iron 1) Ir 2) FE 3) Fe
A. Calcium 2) Ca B. Sulfur 1) S C. Iron 3) Fe
molecules
A molecule is the smallest particle of a pure substance that is capable of a stable independent existence
* Diatomic molecules contain two atoms. • Triatomic molecules contain three atoms. • Polyatomic molecules contain more than three atoms.
Classify the molecules in these diagrams using the terms diatomic, triatomic, or polyatomic molecules. • Solution: H2O2 is a polyatomic molecule, H2O is a triatomic molecule, and O2 is a diatomic molecule
Subtopic
Molecules • Only a few elements exist as individual atoms. • Most elements exist as molecules where two or more atoms of the same element are bonded together. The elements hydrogen, oxygen, phosphorus, and sulfur form molecules consisting of two or more atoms of the same element.
Subtopic
Homoatomic Molecules • The atoms contained in homoatomic molecules areof the same kind. Heteroatomic Molecules • The atoms contained in heteroatomic molecules are of two or more kinds.
Classify the molecules using the terms homoatomic or heteroatomic molecules. • Solution: H2O2 and H2O are heteroatomic molecules and O2 is a homoatomic molecule.
Periodic table
element symbol
Chemical Symbol The symbol that refers to theThe symbol that refers to the elementelement First letter isFirst letter is capitalizedcapitalized, second letter (if, second letter (if applicable) isapplicable) is lowerlower casecase Not all symbols are based on English names forNot all symbols are based on English names for the elements, some come from their Latinthe elements, some come from their Latin names or even other languagesnames or even other languages – Silver – Ag – argentum – Antimony – Sb -stibium – Lead – Pb – plumbum – Copper – Cu – cyprium – Tin – Sn – stannum – Iron – Fe - ferrum – Mercury – Hg - hydrargyrum – Gold – Au - aurum ^
columns
The columns are called FamiliesFamilies oror Groups • Earlier Version had 1-8 followed by A or b Group A elements are called Representative Elements Group B elements are called TransitionTransition ElementsElements • Modern Version labels the columns with 1-18Modern Version labels the columns with 1-18
Summary of Periodic Trends
moving through the periodic table: atomic radius ionization energy electron affinity down group increase decrease less exothermic across a period decrease increase more exothermic
∆U ≡ ionization energy, IE
decreasing size, increasing iE
The minimum energy required to remove an e from the ground state of a gaseous atom • II11 - the energy needed to remove the- the energy needed to remove the first electron from a gaseous atom. K(g) +]+ I1 K+(g) + e–– E = +419 kJ/mol • II22 - the energy needed to remove the second- electron from a gaseous +1 ion. • K+(g) +I2 K2+(g) + e = +3051 kJ/mol ^
decreasing size, increasing IE
history
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Three Types of Elements
• Metals
Shiny when smooth and clean Solid at room temperatur • Only exception - Mercury Good conductors of heat and electricity Most are ductile and malleable
• Metalloids
A.k.a – the semi-metals Boxes bordering the stair-step Physical and chemical characteristics of both metals and non metal
• Nonmetals
Upper Right side of the Periodic Table Generally brittle solids or solids or gases Poor conductors of heat and electricity Bromine is the only liquid at room temperater
The periodic table is useful in predicting:
1. Chemical behavior of the elementsChemical behavior of the elements 2.2. TrendsTrends 3.3. Properties of the elementsProperties of the eleme
Mixtures
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Law of conservation of matter
There is no detectable change in the total quantity of matter present when matter converts from one type to another. • This is true for both chemical and physical changes.
Kinetic Theory of Matte
The kinetic molecular theory of matter is a useful tool for explaining the observed properties of matter in the three different states of solid, liquid and gas. – Postulate 1: Matter is made up of tiny particles called molecules. – Postulate 2: The particles of matter are in constant motion and therefore possess kinetic energy. – Postulate 3: The particles possess potential energy as a result of repelling or attracting each other. – Postulate 4: The average particle speed increases as the temperature increases. – Postulate 5: The particles transfer energy from one to another during collisions in which no net energy is lost from the system.
Particulate Theory of Matter
All matter is made up of tiny particles called molecules and atoms. • Molecules A molecule is the smallest particle of a pure substance that is capable of a stable independent existence. Atoms • Atoms are the particles that make up molecules. Particulate Theory of Matter
Ave. At. Mass = [(% x isotope mass) + (% x isotope mass) + .....]/ total %
sotopes And Atomic Weight Example A specific example of the use of the equation is shown below for the element boron that consists of 19.78% boron-10 with a mass of 10.01 u and 80.22% boron-11 with a mass of 11.01u. this calculated value is seen to agree with the value given in the periodic table.
According to John Dalton, matter is made of very tiny particles called atoms. And atoms cannot be broken down further. • That is atoms were not supposed to be composed of simpler constituents. They were imagined to be like marbles
-All elements are made of tiny atoms. – Atoms cannot be subdivided. – Atoms of the same element are exactly alike. – Atoms of different elements can join to form molecules
3 phases/ states of mater
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isotopes
Isotope Symbols • Isotopes are represented by the symbol AzE • where Z is the atomic number, A is the mass number, and is the elemental symbol.
Isotopes of Carbon and Hydrogen Isotopes of Hydrogen protium deuterium tritium 11H 21H 31 H Isotopes of Carbon 116C 12 6C 136C 146C 156C 166 C
3617Cl OR Cl-36
• Isotopes are atoms that have the same number of protons in the nucleus but different numbers of neutrons. That is, they have the same atomic number but different mass numbers. • Because they have the same number of protons in the nucleus, all isotopes of the same element have the same number of electrons outside the nucleus. ^
Atomic Mass on the Periodic table is the average mass of the isotopes – But the mass number of each isotope is the protons plus the neutrons
atomic mass
Atomic Mass Unit is a unit used to compare the masses of atoms and has the symbol u or amu.
Carbon-12 Chemists have defined the carbon-12 atom as having a mass of 12 atomic mass units. 1 u = 1/12 the mass of a Carbon-12 atom.
the carbon-12 atom
12 atomic mass units.
1 u = 1/12 the mass of a Carbon-12 atom.
1 amu or u is approximately equal to the mass of a single proton or neutron.
Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they have the same atomic number.
matter is anything that has mass and takes up space
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weight is the gravitational force acting on an object and can differ depending on location
scientific notation
a
standerd notation to scientific notation
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Express use scientific calc (EXP) botton 1.8 x 10^-4 in decimal notation. anser 0.00018 Express 4.58 x 10^6 in decimal notation. 4,580,000 On the graphing calculator, scientific notation is done with the button. (EXP) 4.58 x 10^6 is typed 4.58x10 (EXP) 6
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scientific notation is used to express very small or very large numbers and maintain correct number of significant figures
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scientific notation to standard notation
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Subtopic
rules for multiplication
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Rule for Addition and Subtraction
2
Rules for Division
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a x 10 ^n
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Atoms
proton
Credit for the discovery of the proton is debatable pending upon source - However, some people credit Rutherford with its discovery the proton’s mass was eventually identified as 1.672 65 x 10-24 g
discovery of electron
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Atom Is Neutral Atoms have no overall electrical charge so, an atom must have as many electrons as there are protons in its nucleus. Atom -
• Nucleus: Proton + Neutron • Electron ^
discovery of nucleus
Rutherford (1871 – 1973) along with coworkers Geiger, Marsden, and Bohr devised and performed the gold foil experiment in 1911 The experiment: A thin sheet of gold foil was enclosed by a fluor covered circular shroud – There was an opening in the shroud through which alpha particles were shot at the gold foil ^
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The central part of an atom. • Composed of protons and neutrons. • Contains most of an atom's mass
plum pudding model
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Continuous theory of matter – matter could be divided forever without reaching a single smallest indivisible unit
The discontinuous theory of matter – there existed some smallet piece of matter
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neutron
In 1928, a German physicist, Walter Bothe, and his student, Herbert Becker, took the initial step in the search for the neutron. They bombarded beryllium with alpha particles emitted from polonium and found that it gave off a penetrating, electrically neutral radiation, which they interpreted to be high-energy gamma photons.
Eventually, the neutron was discovered in 1932 when James Chadwick used scattering data to calculate the mass of this neutral particle. – Chadwick is credited with his discovery ^
Atoms are the particles that make up molecules.
The Mole Concept Applied To Compounds – The number of molecules in one mole of any compound is called Avogadro's number and is numerically equal to 6.022x1023. – A one-mole sample of any compound will contain the same number of molecules as a one-mole sample of any other compound. – One mole of any compound is a sample of the compound with a mass in grams equal to the molecular weight of the compound.
Examples Of The Mole Concept – 1 mole H2O = 18.02 g H2O = 6.022x1023 H2O molecules – 1 mole CO2 = 44.01 g CO2 = 6.022x1023 CO2 molecules – 1 mole NH3 = 17.03 g NH3 = 6.022x1023 NH3 molecules
The mole concept applied earlier to molecules can be applied to the individual atoms that are contained in the molecules.
CO2
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Examples Of The Mole Concept – 1 mole Na = 22.99 g Na = 6.022x10^23 Na atoms – 1 mole Ca = 40.08 g Ca = 6.022x10^23 Ca atoms – 1 mole S = 32.07 g S = 6.022x10^23 S atoms
Calculate the number of moles of Ca contained in a 15.84 g sample of Ca
15.84 x 1 mole Ca/ 40.08g Ca answer 0.3952 moles Ca
Rutherford – planetary mode
the atom contained a small, positively charged, dense core or center called the nucleus - the electrons traveled around outside the nucleus - most of the atom’s volume was actually empty space - the nuclear diameter is about 10-4 the diameter of the atom
• Criticisms of the planetary model of the atom • Why weren’t the electrons pulled into the positively charged nucleus of the atom? – Rutherford responded by stating the electron’s motion prevents it from being pulled into the nucleus much the same as the planets aren’t pulled into the sun or the moon into the earth • According to classical mechanics (physics) charged particles moving in a curved path should emit energy (light) or some other form of electromagnetic radiation. Eventually, they would lose enough energy to be pulled into the nucleus ^
Subtopic molecules
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Measurements
significant figures
exact number
Exact Numbers do not limit the # of sig figs in the answer
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example
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Average Atomic Mass Formula
sig figs in a measurment include the known digits plus a final estimated dig. sig figs indicate percision of a meassurment
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Count all numbers EXCEPT: • Leading zeros -- 0.0025 • Trailing zeros without a decimal point -- 2,500 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
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Calculating with Sig Figs
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count all numbers exept
leading zeros
4. 0.080 2 sig figs ( first zero is not significant because it is a leading zero but the 0 after 8 is significant
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trailing zeros without a decimal point
examples
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SI units
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Subtopic
Density
Mole
Dimensional analysis/ factor method
A length of rope is measured to be 1834 cm. How many meters is this?
Solution: Write down the known quantity (1834 cm). Set the known quantity equal to the units of the unknown quantity (meters). Use the relationship between cm and m to write a factor (100 cm = 1 m), such that the units of the factor cancel the units of the known quantity (cm) and generate the units of the unknown quantity (m). Do the arithmetic to produce the final numerical answer
1834cm x1m/100cm = 18.34m
The factor- method for solving numerical problems is a four-step systematic approach to problem solving. • Step 1: Write down the known or given quantity. Include both the numerical value and units of the quantity. • Step 2: Leave some working space and set the known quantity equal to the units of the unknown quantity. • Step 3: Multiply the known quantity by one or more factors, such that the units of the factor cancel the units of the known quantity and generate the units of the unknown quantity. • Step 4: After you generate the desired units of the unknown quantity, do the necessary arithmetic to produce the final numerical answer.
From defined quantities: The factors used in the factor-unit method are factors derived from fixed (defined) relationships between quantities. • An example of a definition that provides factors is the relationship between meters and centimeters: 1m = 100cm. This relationship yields two factors: 1m/100cm and 100cm/1m
example
A length of rope is measured to be 1834 cm. How many meters is this? • Solution: Write down the known quantity (1834 cm). Set the known quantity equal to the units of the unknown quantity (meters). Use the relationship between cm and m to write a factor (100 cm = 1 m), such that the units of the factor cancel the units of the known quantity (cm) and generate the units of the unknown quantity (m). Do the arithmetic to produce the final numerical answer. ^
1834cm(1m/100cm)=18.34m
common conversion factors
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example
Example: Convert 34 in. to cm 34 inx 2.54cm/1in= 86cm
Floating topic
Floating topic
Floating topic
electron configuration
– A maximum of 2 electrons per orbital
Exceptional Electron Configurations
Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations. • Exceptions to the aufbau principle are due to
subtle electron-electron interactions in orbitals with very similar energies. • Copper has an electron configuration that is an exception to the aufbau principle. ^
• Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, La, Ce, Gd, Pt, Au, Th, Pa, • U, Np, Cm, Ds, Rg • • A. Half-filled and filled d subshells have unusual stability and this leads to anomalies in electron configurations for some elements • B. Most of these occur with atomic number (Z) > 40, where energy differences between subshells are small. In all cases, the transfer of an electron from one subshell (s) to another (d) lowers the total energy of the atom because of a decrease in electron- electron repulsion. • Niobium Nb*41 [Kr] 4d45s1 • Molybdenum Mo* 42 [Kr] 4d55s1 • Ruthenium Ru*44 [Kr] 4d75s1 • Rhodium Rh* 45 [Kr] 4d85s1 • Palladium Pd* 46 [Kr] 4d105s0 • Silver Ag* 47 [Kr] 4d105s1 • Platinum Pt* 78 [Xe] 4f145d106s0 • Gold Au* 79 [Xe] 4f145d106s! ^
Noble Gas Core Electron Configurations
• Recall, the electron configuration for Na is: Na: 1s2 2s2 2p6 3s1 • We can abbreviate the electron configuration by indicating the innermost electrons with the symbol of the preceding noble gas. • The preceding noble gas with an atomic number less than sodium is neon, Ne. We rewrite the electron configuration: Na: [Ne] 3s1
Condensed Electron Configurations • Neon completes the 2p subshell. • Sodium marks the beginning of a new row. • So, we write the condensed electron configuration for sodium as Na: [Ne] 3s1 • [Ne] represents the electron configuration of neon. • Core electrons: electrons in [Noble Gas]. • Valence electrons: electrons outside of [Noble Gas]. Electron Configurations ^
The periodic table can be used as a guide for electron configurations. • The period number is the value of n. • Groups 1A and 2A have the s-orbital filled. • Groups 3A - 8A have the p-orbital filled. • Groups 3B - 2B have the d-orbital filled. • The lanthanides and actinides have the f-orbital filled. We can use the periodic table to predict which
sublevel is being filled by a particular element.
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Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms.
• We use a numbering system to indicate electrons in an atom • Its given first by a number: 1,2,3,4, etc... dictated by the period number electron shell • Then follows a lower case letter: s,p,d, f... dictated by the group from the periodic table orbital type • Then follows a superscript given over the letter: indicates number of electrons in that orbital
• The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. • The sublevel is written followed by a superscript with the number of electrons in the sublevel. – If the 2p sublevel contains 2 electrons, it is written 2p2
• Aufbau Principle – According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram below, each box represents an atomic orbital. • Pauli Exclusion Principle – According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons. To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired. • Hund’s Rule – Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. ^
• First, determine how many electrons are in the atom. Iron has 26 electrons.(nutral atom electrons and atomic number/#protons is the same) • Arrange the energy sublevels according to increasing energy: –1s 2s 2p 3s 3p 4s 3d ... • Fill each sublevel with electrons until you have used all the electrons in the atom: –Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6 • The sum of the superscripts equals the atomic number of iron (26)
– K+ – As3+
• Write a ground state electron configuration for these ions. O2- Na+ 21
• Write a ground state electron configuration for a neutral atom K Ne
• Write the electron configuration of the neutral atom. • Remove electrons from the orbital with the highest principal QN (value of n) – Fe [Ar] 4s2 3d6 – Fe2+ • A. Fe [Ar] 4s2 3d6 • B. Fe [Ar] 3d6 • C. Fe [Ar] 4s2 3d4 • D. Fe [Ar] 4s1 3d1
An excited atom has an electron or electrons which are not in the lowest energy state. Excited atoms are unstable energetically. The electrons eventually fall to a lower level. * is used to indicate an excited atom. For example: *Li 1s2 3p1. (The ground state for Li is 1s2 2s1.)
• Write an excited electron configuration for the following atoms. • *Al • *K 22
Energy Level, n # of sublevels Letter of sublevels # of orbitals per sublevel# of electrons in each orbital Total electrons inenergy level 1 1 s 1 2 2 2 2 s p 1 3 2 6 8 3 3 s p d 1 3 5 2 6 10 18 4 4 s p d f 1 3 5 7 2 6 10 14 32
The way in which electrons are arranged around the nuclei of atoms.
Arrangement of Electrons in Atoms
SHELLS(n)
total of 7 – 1st energy level is closest to nucleus – Contain Each energy level does not contain the same sublevels – As the distance from the nucleus increases energy levels can hold more electrons –Therefore, they can have more sublevels
If you want to name a specific sublevel in an energy level, you write the energy level number followed by the sublevel. Example: 3rd energy level, d sublevel is written as 3d. ^
SUBSHELLS(i)
ORBITALS (m1)
• Each sublevel contains at least one orbital – Area of higher probability of finding electrons – Every orbital holds 2 electrons – Different sublevels have different shaped orbitals • s = spherical • p = dumbbell
