Download Chapter 6: Section 1 Searching for an Organizing Principle

Chapter 6: Section 1
Searching for an Organizing Principle
Chemists used the properties of elements to sort them into groups
-J.W. Dobereiner published a classification system in which his elements were grouped into triads
-he noted a pattern in his triads- one element in each triad tended to have properties with
values that fell in between the other two
-however, all the known elements could not be grouped into triads
Mendeleev’s Periodic Table
-Dmitri Mendeleev published the first periodic table of elements
-Lothar Meyer published a nearly identical table later that year, but Mendeleev was given more credit
them him because he published his table first and was able to explain its usefulness better
-He wrote the properties of each element on a notecard and moved them around to organize them into
groups based in repeating properties
-Mendeleev arranged the elements in his periodic table in order of increasing atomic mass
-He left blank spaces in his table and predicted that elements would be discovered to fill those spaces
-he even predicted what their properties would be, based on their locations in the table
-gallium and germanium where two elements that he guessed
Periodic Law
-Mendeleev arranged his table by atomic mass
-he didn’t know that the atoms of each element contain a unique number of protons
-this arrangement had problems
-Henry Moseley determined an atomic number for each known element
-In the modern periodic table, elements are arranged in order of increasing atomic number
-There are seven rows or periods
**each period corresponds to a principle energy level
-there are more elements in higher numbered periods because there are more orbitals in higher
energy levels
-Columns = groups
-elements within a group have similar chemical and physical properties
-the pattern of properties within a period repeats as you move from one period to the next
Periodic law: when elements are arranged in order of increasing atomic number, there is a periodic
repetition of their physical and chemical properties
Metals, Nonmetals, and Metalloids
-Each group in the periodic table has 3 possible labels
-red: used in the United States
-blue: used in Europe
-The International Union of Pure and Applied Chemistry (IUPAC) = an organization that sets standards
for chemistry
-they proposed a new system for labeling groups in the periodic table
-they labeled the groups from left to right 1 to 18 in black
-Three classes of elements are metals, nonmetals, and metalloids
-Across a period, the properties of elements become less metallic and more nonmetallic
Metals
-80% of the elements are metals
-good conductors of heat and electric current
-high luster, or sheen caused by the metal‘s ability to reflect light
-solids at room temperature
-ductile (can be pulled)
-malleable (can be bent)
Nonmetals
-these elements are in the upper-right corner of the periodic table
-there is a greater variation in physical properties among nonmetals than among metals
-most are gases at room temperature
-a few are solids
-only one (bromine) is a dark red liquid
-generally, their properties are opposite to those of metals
-poor conductors of heat and electric current
-brittle (shatter easily)
Metalloids
-there is a heavy stair-step line that separates the metals from the nonmetals
-most of the elements that border this line are metalloids
-the line looks like this:
-generally, metalloids have properties that are similar to those of metals and nonmetals
-their behavior can be controlled by changing the conditions
-under some conditions, they can behave like a metal
-under other conditions, they can behave like a nonmetal
Chapter: Section 2
Squares in the Periodic Table
-The periodic table displays the symbols and names of the elements, along with the information about
the structure of their atoms
Atomic number
11
Na
2
8
1
Sodium
22.990
Electrons in each energy level
Element symbol
Element name
Average atomic mass
(See pages 162-163 in the book)
-in the periodic table on pages 162-163, the element symbols are in different colors
-black = solids at room temperature
-red = gases
-blue = liquids (only bromine and mercury)
-green = not found in nature
-Alkali metals = group 1A
-Alkaline earth metals = group 2A
-Halogens = the nonmetals in group 7A
-this includes elements that form salts, particularly chlorine, bromine, and iodine
Electron Configurations in Groups
-electrons play a key role in determining the properties of elements
-elements can be sorted into noble gases, representative elements, transition metals, or inner transition
metals based on their electron configurations
Noble Gases
-includes the elements in group 8A (see periodic table on pg 163)
-these nonmetals are sometimes caused inert metals because they rarely take part in a reaction
-the s and p sublevels are completely filled for these elements
-very stable and unreactive
Representative Elements
-includes elements in groups 1A to 7A
-elements in these groups display a wide range of physical and chemical properties
-includes some metals, some nonmetals, and some metalloids
-most of them are solids, a few are gases, and only bromine is a liquid
-the s and p sublevels of the highest occupied energy level are not filled
-very reactive—tend to gain or lose electrons
-usually unstable
-the elements in group 1A and group 7A= the least stable and most reactive
**for any representative element, its group number = the number of electrons in the highest occupied
energy levels**
-ex: the highest occupied energy level of phosphorus (group 5A) has 5 electrons
Transition Elements
-transition elements = the group B elements (see pages 162-163)
-there are 2 types of transition elements:
1. Transition metals
-the highest occupied s and a nearby d sublevel contain electrons
-these elements are characterized by the presence of electrons in d orbitals
2. Inner transition metals
-the highest occupied s sublevel and an f sublevel contain electrons
-these elements are characterized by the presence of electrons in d orbitals
-they used to be known as rare-earth metals
Blocks of Elements
-the periodic table is divided into sections, or blocks, that correspond to the highest occupied sublevels
-s block = groups 1A and 2A and helium
-p block = groups 3A, 4A, 5A, 6A, 7A, and 8A (not including helium)
-d block = transition metals
-f block = inner transition metals
-These blocks can help you determine electron configurations of elements (see page 166 figure 6.12)
-Each period on the periodic table corresponds to a principle energy level
-ex: sulfur = period 3 and group 6A
-you know that level 1 and 2 orbitals are filled
-you know that there are 4 electrons in the p orbitals by looking at figure 6.12
-so, you can figure out that the electron configuration is: 1s22s22p63s23p4
Chapter 6: Section 3
Trends in Atomic Size
-another way to describe atomic size is to look at molecules (=units that form when atoms of the same
element are joined together)
Atomic radius
molecule
-because the atoms in the molecule are identical, the distance between the two nuclei of the two atoms
can find the size of atoms
-atomic radius = half the distance between the nuclei of two atoms of the same element in a molecule
-this distance is extremely small
-measured in picometers (1012 meters)
-Atomic size increases from top to bottom in a group and decreases from left to right in a period
Group trends in atomic size
-Atomic size increases from top to bottom within groups
-the atomic number increases  the number of filled energy levels increases  atomic radius increases
Why??
-As atomic number increases, the positive charge increases  pulls electrons closer to the nucleus,
making the atomic radius smaller
-However, remember in groups, more energy levels are filled as you go down
-The many occupied energy levels shield electrons from being attracted to the positive nucleus
Periodic trends in atomic size
-Atomic size decreases from left to right across a period
Why??
-more protons  more positive charge
-the more positive charge pulls the electrons in the highest occupied energy level closer to the nucleus
 atomic size decreases
-across a period the electrons are added to the same principal energy level  the shielding effect is
constant
Ions
-ion: an atom with a charge
-atoms are usually electrically neutral with a net charge of zero
-this is because they have equal numbers of protons and electrons
-Positive and negative ions form when electrons are transferred between atoms
-Atoms of metallic elements tend to form cations
-cation = ion with a positive charge
-they tend to lose one or more electrons
-when they lose electrons, there are more protons than electrons, so the ion has a net positive
charge
-ex: Sodium, Na, often becomes Na1+ or Na+ (the one isn’t needed)
-Atomics of nonmetallic elements tend to form anions
-anion = ion with a negative charge
-they tend to gain one or more electrons
-when they gain electrons, there are more electrons than protons, giving the ion a net negative
charge
-ex: Chlorine
Trends in Ionization Energy
-ionization energy = the energy required to remove an electron and overcome the attraction of protons
in the nucleus
-this energy is measured when an element is in its gaseous state
-Low ionization energy = easy to remove an electron
-High ionization energy= hard to remove an electron
-First ionization energy tends to decrease from top to bottom within a group and increase from left to
right across a period
-First ionization energy = the energy needed to remove the first electron from an atom
-the cation produced has a 1+ charge
-Second ionization energy = the energy needed to remove an electron from an ion with a 1+ charge
-the cation produced has a 2+ charge
-Third ionization energy = the energy needed to remove an electron from an ion with a 2+ charge
-the cation produced has a 3+ charge
-Ionization energy can help you predict what ions elements will form
-The increase in energy between the first and second ionization energies large
-It’s pretty easy to remove one electron, but it’s hard to remove a second
Group Trends in Ionization Energy
-Ionization energy decreases from top to bottom within a group
Why??
-Remember: as the size of the atom increases, there are more energy levels that are filled
-The more electron levels are filled electrons are attracted to the nucleus a lot less
-Due to this decreased attraction, less energy is needed for an electron to be removed
Periodic Trends in Ionization Energy
-Ionization energy decreases from top to bottom within a group
Why??
-Nuclear charge increases across the period  the increasing number of protons pulls the electrons in
-The shielding effect stays constant
-Due to this increased attraction, more energy is needed for an electron to be removed
Trends in Ionic Size
-Cations are always smaller than the atoms from which they form
-Anions are always larger than the atoms from which they form
-Metals
-tend to lose electrons and become cations
-the ion is much smaller than the atom
-the attraction of electrons to the nucleus is increased, pulling electrons in and decreasing ionic
size
-the ion has one fewer occupied energy level too (all of the outer electrons were lost)
-Nonmetals
-tend to gain electrons and become anions
-the ion is much larger than the atom
-as the number of electrons increases, the attraction of the electrons to the nucleus decreases
-Periodic trends = ionic size decreases from left to right across a period
-a gradual decrease in the size of cations followed by the gradual decrease in the size of anions
across a period
-Group trends = ionic size increases from top to bottom in groups
Trends in Electronegativity
-electronegativity: the ability of an atom to attract electrons
-can be used to predict the type of bond that will form during a reaction (covalent or ionic)
-scientists use ionization energy to calculate values for electronegativity
-High electronegativity = strong tendency to attract electrons
-tend to gain electrons and form anions
-Low electronegativity = weak tendency to attract electrons
-tend to lose electrons and form cations
-The noble gases aren’t counted – they don’t form compounds or take place in reactions
-Linus Pauling was the first to define electronegativity and won the Nobel Prize
-Paulings are the units used for electronegativity values
-Electronegativity values decrease from top to bottom within a group and increase from left to right
across a period
-metals at the far left of the periodic table = low values
-nonmetals at the far right (excluding noble gases) = high values
-lowest value = cesium
-the least tendency to attract electrons
-it tends to lose electrons and form positive ions when it reacts
-highest value = fluorine
-strongest tendency to attract electrons
- when it reacts, it tends to either attract shared electrons or gain electrons and form anions
Summary of Trends
-The trends that exist among these properties can be explained by variations in atomic structure
-The increase in nuclear charge (which pulls electrons in) explains many trends
-Within groups, an increase in shielding has a very significant effect
(See page 178 figure 6.22—it’s a good summary of the trends)
Stuff Mr. Kub told us in class today that will
probably be on the test
Small size
High ionization energy
High electronegativity
Big size
Low ionization energy
Low electronegativity
Mendeleev = discovered the periodic table
Periodic law = elements in groups have similar properties
Metals
-lustrous, shiny
-malleable
-ductile
-low ionization energy
-conduct electricity well = allows electrons to move around
-electricity = movement of electrons
-especially copper
-conduct heat well
-ex: steel
Nonmetals
-brittle
-poor conductors of heat and electricity
Noble gases
-2 electrons in the s orbital
-6 electrons in the p orbital
-always have fully filled s and p orbitals  full octet  
-they are happy, stable, and nonreactive
Alkali metals
-often lose an electron and then become stable and happy
Electronegativities
-most electronegative = halogens
-the most electronegative element = Fluorine
-least = alkali metals
-the least electronegative element = Cesium
Cations = smaller than neutral atoms
Anions = larger than neutral atoms
Remember:
s = 1 orbital
p = 3 orbitals
d = 5 orbitals
f = 7 orbitals