powerpoint presentation about elements

The Periodic Table and Periodic Law �

Development of the Modern Periodic Table �

Aristotle �

• Four element theory: earth, air, fire & water

Antoine Lavoisier �

• Wrote the first extensive list of elements containing 23 elements.�Distinguished between metals and non-metals.
• Some of Lavoisier's elements were later shown to be compounds and mixtures.

Industrial Revolution �

• Great increase in number of known elements due to electricity and spectrometry.
• Needed an organized system for classification.
• New industries – soaps, dyes, fertilizers, petrochemicals…
• Increase in chemical pollution.
• Chemists agree on method for determining atomic masses.

John Newlands �

• The known elements (>60) were arranged in order of atomic weights and observed similarities between the first and ninth elements, the second and tenth elements, etc. He proposed the 'Law of Octaves'.
• Newlands' Law of Octaves identified many similarities amongst the elements, but also required similarities where none existed. He did not leave spaces for elements as yet undiscovered. Forerunner to the notion of periods.

Lothar Meyer �

• Compiled a Periodic Table of 56 elements based on the periodicity of properties such as molar volume when arranged in order of atomic weight.
• Meyer & Mendeleev produced their Periodic Tables simultaneously.

Dmitri Mendeleev �

• Published a table based on atomic weights but arranged 'periodically' with elements with similar properties under each other.
• Mendeleev's Periodic Table was important because it enabled the properties of elements to be predicted by means of the ' P eriodic law ': properties of the elements vary periodically with their atomic weights.

Mendeleev's Periodic Table

• made a table in which he arranged the elements in order of increasing atomic mass into columns with similar properties.
• Gaps were left for elements that were unknown at that time and predicted the properties of three elements that were undiscovered at the time. (the elements were gallium, scandium and germanium).
• When these elements—scandium, gallium, & g ermanium—were soon discovered and found to have most of the predicted properties
• The major changes in the organization of the periodic table since Mendeleev’s time
• New elements have been added
• elements are now arranged in order of atomic number instead of atomic mass.

William Ramsay

• Discovered the Noble Gases.
• Ramsay removed oxygen, nitrogen, water and carbon dioxide from a sample of air and was left with a gas 19 times heavier than hydrogen, very unreactive - Argon. In 1895 he discovered helium as a decay product of uranium and matched it to the emission spectrum of an unknown element in the sun. (helios is Greek for Sun). He later discovered neon, krypton and xenon. Ramsay was awarded a Nobel Prize in 1904.

Henry Moseley

• Determined the atomic number of each of the elements.
• He modified the Periodic Law to read that the properties of the elements vary periodically with their atomic numbers.
• Moseley's modified Periodic Law puts the elements tellurium and iodine in the right order, as it does for argon and potassium, cobalt and nickel.

Glenn Seaborg �

• Synthesised transuranic elements (the elements after uranium in the periodic table)
• In 1940 uranium was bombarded with neutrons in a cyclotron to produced neptuniun (Z=93). Plutonium (Z=94) was produced from uranium and deuterium. These new elements were part of a new block of the Periodic table called Actinides. Seaborg was awarded a Nobel Prize in 1951.

The modern periodic table

• Study the modern periodic table that appears in your textbook on pages 178–179.
• Each square gives certain information about each element, as shown in the following diagram.

Periodic Table Concept Map

• The elements are arranged in the periodic table in order of increasing atomic number in horizontal rows called periods.
• Because the pattern of properties repeats in each new row of elements, the elements in a column have similar properties and are called a group or family of elements. The groups are designated with a number.
• Groups 1 and 2 and 13–18 are called the main group or representative elements.
• The elements in groups 3–12 are called the transition elements.

3 main classes

• Elements are divided into three main classes—metals, metalloids, and nonmetals.
• As you can see from the periodic table, the majority of the elements are metals. Metals are generally shiny solids and are good conductors of heat and electrical current. Some groups of elements have names. For example, the first two groups of metals, groups 1 and 2, are called the alkali metals and the alkaline earth metals, respectively.
• Most of the elements to the right of the heavy stair-step line in the periodic table are nonmetals, which are generally either gases or brittle solids at room temperature. Group 17 elements are commonly called halogens, and group 18 elements are the noble gases.
• Many of the elements that border the stair-step line are metalloids, which have some of the characteristics of both metals and nonmetals.
• A groups = representative elements
• B groups = transition metals
• Shiny, solid at room temperature.
• Good conductors of heat and electricity.
• Malleable- can be pounded into thin sheets.
• Ductile – can be drawn into wires.
• Generally gases or brittle solids.
• Poor conductors of heat and electricity.
• Exception: Bromine is a liquid at room temperature.
• Most of the nonmetals are on the far right of the periodic table.
• Semimetals with properties of both metals and nonmetals.
• Stair step between metals and nonmetals on periodic table.
• Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium, Astatine

Transitions Metals

• B Group elements that include Lanthanide and Actinide (at the bottom of the periodic table)
• Lanthanides act as phosphors that emit light when struck by electrons.

Group 1: Alkali Metals

• Valence # = 1
• Charge = +1
• Characteristics:
• Lose one valence electron (s 1 ) and form a 1+ ion
• Soft enough to cut with a knife
• Highly reactive; found in nature combined with other elements
• Good conductors of heat and electricity

Group 2: Alkaline Earth Metals

• Valence# = 2
• Charge = +2
• Produce alkaline solutions when mixed with water, except for Be
• Shiny solids harder than alkali metals
• Usually found combined with other elements
• Lose 2 valence electrons (s 2 ) and form 2+ ions

Group 3-12: Transition Metals

• Valence # = 1 or 2
• Electrical conductivity, luster and malleability
• Hard solids with high melting and boiling points
• Lanthanides are silvery metals with high melting points; found mixed together in nature
• Actinides are radioactive and only 3 exist in nature; the rest are synthetic transuranium elements

Group 13: Boron

• Valence # = 3
• Charge = 3+
• Valence electrons are s 2 p 1
• Always found combined with other elements in nature
• B, Al, Ga, & In form 3+ ions

Group 14: Carbon

• Valence # = 4
• Charge = 4+/-
• Valence electrons are s 2 p 2
• The group has one nonmetal, two metalloids and two metals
• Carbon is not representative of the other members of the group

Group 15: Nitrogen

• Valence # = 5
• Charge = 3-
• Valence electrons are s 2 p 3
• Wide varieties of properties, for example N forms a 3- ion and Bi forms a 3+ ion

Group 16: Oxygen

• Valence = 6
• Charge = 2-
• Six valence electrons (s 2 p 4 ) and members act mainly as nonmetals
• Members tend to gain 2 electrons and form 2- ions
• Members can also share 2 electrons to achieve stability

Group 17: Halogen

• Valence # = 7
• Charge = 1-
• Form compounds with almost all metals; except astatine
• Share one or lose one electron to become stable
• Valence electrons are s 2 p 5 ; form 1- ions

Group 18: Noble Gases

• Full outer shell, 8 valence electrons
• Colorless and unreactive
• Stable electron configurations; valence electrons: He s 2 , all others s 2 p 6

Classification of the Elements �

• Chemists today understand that the repetition of properties of elements occurs because the electron configurations of atoms exhibit repeating patterns. Thus, the arrangement of elements in the periodic table reflects the electron structures of atoms. For example, the group number of a representative element gives the number of valence electrons in an atom of that element. The group 13 element aluminum, for example, has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 1 , or [Ne]3s 2 3p 1 . The three electrons in the third energy level (3s 2 and 3p 1 ) are the valence electrons of the aluminum atom. In a similar way, the period number of a representative element indicates the energy level of the valence electrons. Aluminum is in the third period, and aluminum’s valence electrons are in the third energy level.

Blocks of elements

• The periodic table is divided into blocks of elements that correspond to the energy sublevel being filled as you move across a period. The 1 and 2 groups constitute the s-block elements because their highest-energy electrons are in s orbitals. The remaining groups of representative elements, 13 through 18, make up the p-block of elements. In these elements, s orbitals are filled, and the highest-energy electrons are in p orbitals. The transition metals are the d-block elements. In these elements, the highest-energy electrons are in the d sublevel of the energy level one less than the period number. Most d-block elements have two electrons in s orbitals, but in some, such as chromium and copper, the d sublevel “borrows” an electron from the s orbital to form half-filled (Cr) or filled (Cu) d orbitals. The remaining block is the f-block, or inner transition metals. The highest-energy electrons in these elements are in an f sublevel of the energy level two less than the period number. The following Example Problem illustrates how electron configuration determines an element’s position in the periodic table.

Example Problem

• Electron Configuration and the Periodic Table
• The electron configuration of phosphorus is [Ne]3s 2 3p 3 . Without using the periodic table, determine the group, period, and block in which the element is located in the periodic table. First, identify the valence electrons and note their energy level. In phosphorus, the 3 in front of the s and p orbitals indicates that the valence electrons are in the third energy level. Therefore, phosphorus will be found in the third period of the periodic table. Next, note the sublevel of the highest-energy electrons. In the case of phosphorus, these electrons are in a p sublevel. Therefore, phosphorus will be found in the p-block. Finally, use the number of valence electrons to determine the group number of the element. Phosphorus has two electrons in an s orbital and three electrons in p orbitals for a total of five valence electrons. Because there are no incomplete d or f sublevels, phosphorus must be a representative element in group 15.
• To summarize, phosphorus is found in period 3, group 15, and the p-block of the periodic table. A glance at the table will confirm this answer.
• Coulomb’s Law – a basic law of physics that tells us the FORCE between charged particles depends on the charge and the distance between the charges
• Shielding – valence electrons are repulsed by the inner core electrons; results is that valence electrons are screened from the full charge of the nucleus
• Effective Nuclear Charge - The only part of the nuclear charge that valence electrons actually get

Z eff = # of protons - # of inner core electrons

Coulomb’s Law

• q A and q B are the charges
• q A would be the charge (-1) of valcence electron
• q B is the charge of the nucleus (changes with every element
• d is the distance between nucleus and a valence electron
• k is Coulomb’s constant; we do not need this to understand periodic trends

Practice with Coulomb’s Law

• How would the force change if the value charge on the nucleus doubles?
• How would the force change if the distance between the nucleus and the valence electron doubled?

Periodic Trends

• The arrangement of the periodic table was developed around more elemental characteristics than atomic number.
• Properties that tend to change in a predictable way are called TRENDS.

There are many different Periodic trends! We are going to look at Atomic Radius, Ionization Energy, and Electronegativity.

Periodic Trends �

• The electron structure of an atom determines many of its chemical and physical properties. Because the periodic table reflects the electron configurations of the elements, the table also reveals trends in the elements’ chemical and physical properties.

Definitions

• Atomic Radius – refers to the size of an atom; half the distance between adjacent nuclei;
• Ionic Radius – size of ion; half distance between adjacent nuclei;
• Cations : smaller than the original atom
• Anions : larger than the original atom
• Ionization Energy – minimum energy required to remove an electron from an atom (gas state)
• Electron Affinity – the energy change that occurs when an electron is added to a gaseous atom
• Electronegativity – an atom’s attraction for electrons in a chemical bond

Atomic radius

• The atomic radius is a measure of the size of an atom. The larger the radius, the larger is the atom. Research shows that atoms tend to decrease in size across a period because the nuclei are increasing in positive charge while electrons are being added to sublevels that are very close in energy. As a result, the increased nuclear charge pulls the outermost electrons closer to the nucleus, making the atom smaller. Moving down through a group, atomic radii increase. Even though the positive charge of the nucleus increases, each successive element has electrons in the next higher energy level. Electrons in these higher energy levels are located farther from the nucleus than those in lower energy levels. The increased size of higher energy level outweighs the increased nuclear charge. Therefore, the atoms increase in size.

Explain Atomic Radius

• Why does the atomic radius increase as you go down a column?
• Why does the atomic radius decrease as you go left to right across a row?

ANS: As you go down the table, there are more energy levels (rings).

The valence electrons are located further from the nucleus

ANS: All atoms on a row of the periodic table have their valence

electrons on the same energy level. In other words, they are

the same distance from the nucleus. However, each element

adds another proton to the nucleus. The inner core does not change.

Therefore, the effective nuclear charge increases. According to

Coulomb’s Law , this increase in charge will increase the force.

Inner Core Same

For both elements

More protons –

Greater Nuclear

Ionic radius

• When an atom gains or loses one or more electrons, it becomes an ion. Because an electron has a negative charge, gaining electrons produces a negatively charged ion, whereas losing electrons produces a positively charged ion. As you might expect, the loss of electrons produces a positive ion with a radius that is smaller than that of the parent atom. Conversely, when an atom gains electrons, the resulting negative ion is larger than the parent atom. Practically all of the elements to the left of group 14 of the periodic table commonly form positive ions. As with neutral atoms, positive ions become smaller moving across a period and become larger moving down through a group. Most elements to the right of group 14 (with the exception of the noble gases in group 18) form negative ions. These ions, although considerably larger than the positive ions to the left, also decrease in size moving across a period. Like the positive ions, the negative ions increase in size moving down through a group.

Ionic Radius

An isoelectronic series is a sequence of species all having the same number of electrons (and thus the same amount of electron-electron repulsion) but differing in nuclear charge. Of course, only one member of such a sequence can be a neutral atom (neon in the series shown below.) The effect of increasing nuclear charge on the radius is clearly seen.

Ionization energy

• Energy is required to pull an electron away from an atom. The first ionization energy of an element is the amount of energy required to pull the first valence electron away from an atom of the element. Atoms with high ionization energies, such as fluorine, oxygen, and chlorine, are found on the right side of the periodic table and are unlikely to form positive ions by losing electrons. Instead, they usually gain electrons, forming negative ions. Atoms with low ionization energies, such as sodium, potassium, and strontium, lose electrons easily to form positive ions and are on the left side of the periodic table. Recall that atoms decrease in size from left to right across a period. First ionization energies generally increase across a period of elements primarily because the electrons to be removed are successively closer to the nucleus. First ionization energies decrease moving down through a group of elements because the sizes of the atoms increase and the electrons to be removed are farther from the nucleus.

Ionization Energy Chart

Ionization Energy continued�

Note the very large jumps in the energies required to remove electrons from the 1s orbitals of atoms of the second-row elements Li-Ne.

Explain Ionization Trends

• Why does the ionization energy decrease as you go down a column?
• Why does the ionization energy increase as you go from left to right across a row?

Ans: As you go down a column, the valence electrons are on levels

that are further away from the nucleus. According to Coulomb’s

law, the further the distance, the less force holding the electrons, so

they are easier to lose.

Ans: Across a row, the nuclear charge increases while the shielding

effect of the inner core remains the same. According to Coulomb’s

law, increasing nuclear charge creates a greater force, so the valence

electrons are held more tightly.

Electron Affinity Chart

The octet rule

• When atoms lose or gain electrons, they generally do so until the ion has eight valence electrons—the stable s 2 p 6 electron configuration of a noble gas. This principle is called the octet rule. Exceptions to this rule are hydrogen, which can gain an electron, obtaining the stable 1s 2 configuration of helium, and elements in period 2, such as lithium and beryllium that lose electrons, also obtaining the helium configuration.
• The octet rule lets you predict the ionic charge of a representative element. For example, you can predict that an element in group 16, having high ionization energy, will gain two electrons to achieve a stable octet configuration.

Electronegativity

• When atoms combine chemically with each other, they do so by forming a chemical bond. This bond involves either the transfer of electrons or sharing of electrons to varying degrees. The nature of the bond between two atoms depends on the relative ability of each atom to attract electrons from the other, a property known as electronegativity. The maximum electronegativity value is 3.98 for fluorine, the element that attracts electrons most strongly in a chemical bond. The trends in electronegativity in the periodic table are generally similar to the trends in ionization energy. The lowest electronegativity values occur among the elements in the lower left of the periodic table. These atoms, such as cesium, rubidium, and barium, are large and have few valence electrons, which they lose easily. Therefore, they have little attraction for electrons when forming a bond. Elements with the highest electronegativity values, such as fluorine, chlorine, and oxygen, are found in the upper right of the periodic table (excluding, of course, the noble gases, which do not normally form chemical bonds). These atoms are small and can gain only one or two electrons to have a stable noble-gas configuration. Therefore, when these elements form a chemical bond, their attraction for electrons is large. Electronegativity generally increases across a period and decrease down through a group.

Summary Table

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The Periodic Table of Elements

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Elements and Compounds

- Elements are pure substances that cannot be broken down further without losing their identity, and there are currently 118 known elements, with 88 occurring naturally. - In the universe, hydrogen makes up 75% and helium 20%, while on Earth oxygen is the most abundant element in the crust at 46.6% and silicon is second most at 27.7%. - Compounds are pure substances made of two or more chemically bonded elements, with properties different from the individual elements, and can be represented by chemical formulas showing the elements and their ratios. Read less

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• 3. Elements There are currently 118 elements that have been identified though only 88 of them are naturally occurring.
• 4. Elements in the Universe In our universe, hydrogen makes up 75% of all matter! Helium makes up about 20% with oxygen being the 3 rd most abundant element. All of the other elements are relatively rare in the universe.
• 5. Elements in the Earth In the Earth’s crust, oxygen is the most abundant element (46.6%). Silicon is the second most abundant element (27.7%). Aluminum (8.1%), iron (5.0%), calcium (3.6%), sodium (2.8%), potassium (2.6%). and magnesium (2.1%) complete the list of elements that account for approximately 98.5% of the total mass of the earth's crust.
• 6. Elements Elements are a pure substance. Made of only one kind of material, has definite properties, and is the same all throughout. Elements are the simplest pure substance. They cannot be broken down into simpler substances without losing their identity.
• 7. Elements and Atoms The smallest particle of an element that has the properties of that element is called an atom . Atoms: the building blocks of matter. Atoms of the same element are alike; atoms of different elements are different.
• 8. Chemical Symbols Shorthand way of representing the elements. Usually one or two letters. Usually taken from the name of the element. Carbon-C, Calcium-Ca, Hydrogen-H, Iodine-I, Oxygen-O, Chlorine-Cl
• 9. Chemical Symbols Some symbols come from their Latin name: Gold-Au--aurum Silver-Ag--argentum Iron-Fe--Ferrum Mercury-Hg--hydrogyrum
• 10. Compounds Pure substances made up of more than one element. 2 or more elements chemically combined. Ex: H 2 O, NaCl, C 6 H 12 O 6 , CO 2 Unlike elements, compounds can be broken down to simpler substances. This can happen through a chemical reaction.
• 11. Compounds The properties of the elements that make up a compound are often quite different from the properties of the compound itself. Sodium-Na--highly reactive metal Chlorine-Cl--poisonous gas Sodium Chloride-NaCl--table salt
• 12. Molecules Compounds are made of molecules. A molecule is 2 or more atoms chemically bonded. Water-2 atoms of hydrogen and one atom of oxygen-together they form one molecule of H 2 O. A molecule is the smallest particle of a compound that has all the properties of that compound. Just as all atoms of a certain element are alike, all molecules of a certain compound are alike.
• 13. Chemical Formulas A shorthand way of representing compounds. If chemical symbols are the “letters,” these are the “words.” Ex: NH 3 - ammonia, C 3 H 7 OH - rubbing alcohol Sometimes, the formula represents a molecule of a single element. These are called diatomic molecules. This is how that element is naturally found. O 2 -Oxygen H 2 -Hydrogen Cl 2 -Chlorine
• 14. Chemical Formulas Subscripts are small numbers used in chemical formulas. They are placed to the lower RIGHT of the chemical symbols. Represent # of atoms of an element in a compound. CO 2 = 1 atom of carbon and 2 atoms of oxygen. H 2 SO 4 = 2 atoms of hydrogen, 1 atom of sulfur and 4 atoms of oxygen