unit 2 – chapters 2, 8, & 9
DESCRIPTION
Unit 2 – Chapters 2, 8, & 9. The Components of Matter. Definitions for Components of Matter. - PowerPoint PPT PresentationTRANSCRIPT
Unit 2 – Chapters 2, 8, & 9
The Components of Matter
Definitions for Components of Matter
________________ - the simplest type of substance with unique
physical and chemical properties. An element consists of only
one type of atom. It cannot be broken down into any simpler
substances by physical or chemical means.
__________________- a structure that
consists of two or more atoms that are
chemically bound together and thus behaves
as an independent unit.
__________________ - a substance
composed of two or more elements
which are chemically combined.
________________ - a group of two or more elements and/or compounds that are physically intermingled.
Definitions for Components of Matter
• __________________ (1766-1844), an English schoolteacher and chemist, studied the results of experiments by Lavoisier, Proust, and many other scientists.
• Dalton proposed his atomic theory of matter in 1803.
• Although his theory has been modified slightly to accommodate new discoveries, Dalton’s theory was so insightful that it has remained essentially intact up to the present time.
Dalton’s Atomic Theory
Dalton’s Atomic TheoryDalton’s Atomic Theory
1. All matter consists of ______________________.
2. Atoms of one element ___________________ be converted into atoms of another element.
3. Atoms of an element are _____________________ in mass and other properties and are _______________________ from atoms of any other element.
4. ______________________________ result from the chemical combination of a specific ratio of atoms of different elements.
The Postulates
Structure of the Atom
• J.J. (John Joseph) Thomson, physicist• 1890-1900• Showed that the atoms of any element can be
made to emit tiny negative particles - called _______________________.
• Thompson knew that the entire atom was not negatively charged so he concluded that the atom must also contain positive particles that balance the negative charge, giving the atom a ___________________________________.
• Ernest Rutherford• 1911• Learned physics in J.J. Thomson’s
laboratory in the late 1890s.• Main area of interest was the
_____________________________ - positively charged particles with a mass approximately 7500 times that of an electron.
Ernest Rutherford
• By 1919, Rutherford concluded that the nucleus of an atom contained what he called ______________________(has the same magnitude of charge as the electron, but its charge is positive)
• Protons have a _________ charge and the electron a charge of ______________.
• 1932, he and a coworker (James Chadwick) were able to show that most nuclei also contain a neutral particle that they named the _________________ (which has no charge)
Modern Concept of Atomic Structure
• The simplest view of the atom is that it consists of a tiny nucleus that is about 10-13 cm in diameter.
• Electrons move about the nucleus at an average distance of about 10-8 cm from it.
• Nucleus contains ___________________, which have a positive charge equal in magnitude to the electron’s negative charge, and ________________________, which have almost the same mass as a proton but no charge.
Modern Concept of Atomic Structure
• Mass and charge of the electron (e-), proton (p+), and neutron (N)
The mass and charge of the electron, proton, and neutron.
Particle Relative Mass* Relative ChargeElectron 1 1-Proton 1836 1+Neutron 1839 None*The electron is assigned a mass of 1 for comparison
Distinguishing Between Atoms
• Protons and electrons are _______________ in an atom of an element (neutral charge).
• The ________________________ of an element is the number of ______________ in the nucleus of an atom of that element. (If the p+ and e- are the same, then the atomic number will also identify the number of e-)
Distinguishing Between Atoms
• The sum of the number of neutrons and the number of protons in a given nucleus is called the atom’s ____________________________.
protons + neutrons = mass number
• _____________________• atoms with the same number of protons but different
numbers of neutrons.• Elements on the periodic table are the most common
isotopes of those substances.
Distinguishing Between Atoms
• Isotopes• Because they have different numbers of
_________________________, their mass numbers will be different.
• Neon - 20• Neon - 21• Neon - 22• All of these are isotopes of neon.
Distinguishing Between Atoms
• Isotopes• 3 known isotopes of hydrogen
• hydrogen - 1 [hydrogen]• hydrogen - 2 [deuterium]• hydrogen - 3 [tritium]
Isotopic Symbols
• X = the symbol of the element• A = the mass number • Z = the atomic number
A
Z X
1
1 H
Hydrogen
2
1 H
Deuterium
3
1 H
Tritium
Atomic Masses
• Because atoms are so tiny, the normal units of mass - the gram and the kilogram - are much too large to be convenient.
• Mass of a single carbon atom is 1.00 x 10-23 grams.• When describing the mass of an atom, scientists
have defined a much smaller unit of mass called the __________________________.
Atomic Masses
• In terms of grams:• 1 amu = atomic weight of a substance
expressed in grams• 1 carbon atom = 12.01 amu = 12.01 grams• 1 aluminum atom = 26.98 amu = 26.98
grams
Periodic Table of Elements
Periodic Table of Elements
• Shows all the known elements and gives a lot of information about each element.
• Invaluable in chemistry!
Development of Periodic Table
• Elements in the same group generally have similar ________________________________.
• Properties are not identical, however.
Development of Periodic Table
Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.
Development of Periodic Table
Mendeleev, for instance, predicted the discovery of germanium as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.
Dmitri Mendeleev
• Organized the elements according to their increasing __________________________.
• Then he grouped them into columns and rows according to physical and chemical properties.• Row – __________________• Column - __________________
Henry Moseley
• Rearranged the elements according to their __________________________.
• Arranging the elements in this manner provided for a better fit of chemical and physical properties and aligned those elements that were discovered after Mendeleev developed the original periodic table.
Parts of the Periodic Table of Elements
• _____________________ – substances to the left of the dark line
• _____________________ – substances to the right of the dark line
• _____________________ – those elements that border the line
Properties of Metal, Nonmetals,and Metalloids
Metals versus Nonmetals
Differences between metals and nonmetals tend to revolve around these properties.
Metals versus Nonmetals
• Metals tend to form ____________.• Nonmetals tend to form _____________.
Metals
Tend to be ______________, ______________, ______________, and good conductors of ______________ and ______________.
Metals
• Compounds formed between metals and nonmetals tend to be ______________.
• Metal oxides tend to be ______________.
Nonmetals
• ______________, ______________substances that are ______________ conductors of heat and electricity.
• Tend to gain ______________ in reactions with metals to acquire noble gas configuration.
Nonmetals
• Substances containing only nonmetals are ______________ compounds.
• Most nonmetal oxides are ______________.
Metalloids
• Have some characteristics of ______________, some of ______________.
• For instance, silicon looks shiny, but is brittle and fairly poor conductor.
Fireworks• Potassium – combustible element that helps oxidize firework
mixtures• Lithium – adds red color • Sodium – gold and yellow colors• Magnesium – bright white color• Calcium – deepens the colors of the other elements in the
fireworks• Strontium – red color and stablizes other elements• Barium – green color and stablizes other elements• Titanium – produces the spark• Iron – produces sparks• Copper – blue color• Zinc – smoke clouds• Aluminum – silver and white sparks and flames – sparklers• Carbon – black powder• Phosphorus – fuel • Sulfur – fuel• Antimony – glitter effects
Groups of the Periodic Table
• Group 1 – Alkali metals• Group 2 – Alkaline Earth metals• Group 11 – Coinage metals• Group 17 – Halogens• Group 18 – Noble Gases
Electronic Structure
Electronic Structure
• ______________ are vital in the determining the properties of elements.
• Electrons are involved in bonding between ______________.
• Electrons move around the nucleus in different ____________________________.
Energy Levels
• There are ______________ energy levels – one for each corresponding row or period on the periodic table.
• Within each energy level there are ______________.
Sublevels
• Within each energy level there are a possibility of 4 sublevels – depending on which energy level you are dealing with.
Sublevels
• The 4 different sublevels are:
• _____ – holds a maximum of 2 electrons• _____ – holds a maximum of 6 electrons• _____ – holds a maximum of 10 electrons• _____ – holds a maximum of 14 electrons
Sublevels
• Energy level 1 – only has an “s” sublevel• Energy level 2 – only has an “s” and “p”
sublevel• Energy level 3 – only has an “s”, “p”, and “d”
sublevel• Energy level 4 & 5 – have an “s”, “p”, “d”, and
“f” sublevels• Energy level 6 – only has an “s”, “p”, and “d”
sublevel• Energy level 7 – only has an “s” and “p”
sublevel
Sublevels
• Level 1 – maximum of ______________• Level 2 – maximum of ______________• Level 3 – maximum of ______________• Level 4 & 5 – maximum of ___________• Level 6 – maximum of ______________• Level 7 – maximum of ______________
Orbitals
• Within each sublevel, there are orbitals – locations where the electrons are actually located.
• An orbital can hold 2 electrons only.• Sublevel “s” – _____orbital (total of 2 e-)• Sublevel “p” – _____orbitals (total of 6 e-)• Sublevel “d” – _____orbitals (total of 10 e-)• Sublevel “f” – _____orbitals (total of 14 e-)
Electron Configuration
• Shows the location of the electrons in an atom of an element.
• ______________– electrons fill specific energy levels and sublevels from the nucleus out in a specific order.
Electron Configuration
• 1. Determine the number of electrons an atom of an element has.
• 2. Fill the energy levels and sublevels in order using the diagonal rule.
• 3. Write the configuration
Electron Configuration
• Sodium – 11 electrons
• 1s2 2s2 2p6 3s1
• Follow the diagonal for filling the energy levels and sublevels
• If the sublevel is filled, include all the electrons, however, when only one is needed (as in the 3s1), only include one electron.
• If you total the superscript numbers (number of electrons), there are only 11, which is the number of electrons that sodium has.
Electron Configuration
• Bromine – 35 e-
• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
• Notice that after the 4s it is 3d. You must go in the specific order as shown by the diagonal rule.
Orbital Notation
• ______________shows the electrons within each energy level and sublevel and how they ______________.
• We know that 2 ______________ can exist within an orbital.
• We also know that electrons are all ______________ charged.
• For the 2 electrons to exist in an orbital they must spin in opposite directions.
Orbital Notation
• ______________– electrons will remain as unpaired as possible.• Minimizes
____________________________– everyone gets their own room whenever possible.
• Pauli’s Exclusion Principle• Electrons will spin in opposite directions
e- Configuration & Orbital Notation
__ __ __ __ __ __• 1s2 2s2 2p6 3s1
• The lines above show the maximum number of orbitals for each sublevel.
• All possible orbitals must be shown even if electrons do not exist in the orbital.
e- Configuration & Orbital Notation
__ __ __ __ __ __1s2 2s2 2p6 3s1
• Notice that the arrows go in opposite directions – shows the electrons spinning in opposite directions.
• The last energy level / sublevel only has one arrow because there is only one electron in that sublevel.
• http://intro.chem.okstate.edu/APnew/Default.html
• Website with simulation
Valence Electrons
• ______________ electrons are the electrons in the highest energy level, “s” and “p” sublevel.• Totaling 8 electrons – 2 in the “s” and 6 in
the “p”
• These are the electrons involved in ______________.
Quantum Numbers
The Uncertainty Principle
• Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known.
• In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!
• We can only calculate the probability of finding an electron within a given space.
Quantum Mechanics
• Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.
• It is known as _________________.
Quantum Numbers
• Each orbital describes a spatial distribution of electron density.
• An orbital is described by a set of three ___________________.
Principal Quantum Number, n
• The principal quantum number, n, describes the ____________________ on which the orbital resides.
• The values of n are integers of numbers ranging from 0 to infinity.
• The larger the n value, the larger the ______________.
• As n increases, the energy increases. • Lower energy = more __________ atom (n
would be small)
Principal Quantum Number, n
• An electron can jump from a lower energy state to another by emitting or absorbing ______________.• Absorb energy – electrons can move to a
higher energy state.
Azimuthal Quantum Number, l
• This quantum number defines the _______________ of the orbital.
• We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals.
• 0 to n-1
Azimuthal Quantum Number, l
Value of l 0 1 2 3
Type of orbital s p d f
Magnetic Quantum Number, ml• Describes the three-dimensional
orientation of the orbital.• Values are integers ranging from -l to l• Assign the “blanks” in orbital notation
with zero on the middle blank and then –l through zero to +l.
Magnetic Quantum Number, ml
• Orbitals with the same value of n form a shell.• Different orbital types within a shell are
subshells.
s Orbitals
• Value of l = 0.• Spherical in shape.• Radius of sphere
increases with increasing value of n.
p Orbitals
• Value of l = 1.• Have two lobes with a node between them.
d Orbitals
• Value of l is 2.
• Four of the five orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.
Spin Quantum Number, ms
• In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy.
• The “spin” of an electron describes its _________________, which affects its energy.
Spin Quantum Number, ms
• This led to a fourth quantum number, the spin quantum number, ms.
• The spin quantum number has only 2 allowed values: +½ and -½.
Spin Quantum Number, ms
• Electrons spin on an ____________• Electron spin quantum number• Electrons in the same orbital spin in
opposite directions• Values of +½ and –½
Pauli Exclusion Principle
• No two electrons in the same atom can have exactly the same energy.
• For example, no two electrons in the same atom can have identical sets of quantum numbers.
• The first electron placed in an orbital gets the +½ and the second one gets the -½.
Sulfur• Electron Configuration for sulfur is:
• 1s2 2s2p6 3s2p4
• Highest energy level is 3, so n = 3• The last electron is in the p orbital, so l = 2
• ___ ___ ___
• The last electron is in the -1 position, so ml = -1
• That electron is the 2nd electron in that orbital, so it has an ms = - ½
-1 0 +1
Quantum Numbers for Sulfur
• 3, 2, -1, - ½