Periodic Table Elements Bonding

Periodic Table/Elements/Bonding (Groups and Families) ➔ The Periodic Table is an array of all 118 chemical elements ➔ As mentioned before, a chemical element is a chemical substance that cannot be broken down by chemical means ➔ The elements are arranged in increasing atomic number ◆ The atomic number is the number of protons in an atom of the specific element ◆ It will always be a whole number ➔ To the right, you will see the element “Hydrogen.” Above its symbol “H,” there is a number. That number is Hydrogen’s atomic number. ➔ Therefore, the number above an element’s symbol is its atomic number ➔ This means an atom of Hydrogen only has one proton ➔ In a neutral atom, the number of protons is equal to the number of electrons ➔ Therefore, a neutral atom of hydrogen would have one electron since it has one proton according to the atomic number ➔ The atomic number determines the element. There is only one element with an atomic number of one: Hydrogen As you move through the periodic table, the atomic number of elements increases by one. ◆ The second element is Helium. Thus, it has an atomic number of two, meaning an atom of Helium contains two protons ◆ According to this picture of Neon on the periodic table, what is Neon’s atomic number? ◆ Answer: 10 (because “10” is above the symbol “Ne”) ◆ This means it is the tenth element on the periodic table and each atom of Neon contains ten protons ➔ Symbol of Chemical Element ➔ The symbol is the abbreviated and shortened version of the element’s name, allowing each symbol to be easily identified on the periodic table ◆ The symbol is either one or two letters, and it comes from the Latin name of the element ◆ If there is one letter, it must be capitalized ◆ If there are two letters, the first one is capitalized and the second one is lowercase ● For example, Helium’s chemical symbol is “He” ● The chemical symbol for Boron is “B” ● As shown in the picture, Hydrogen’s symbol is “H” ➔ The mass number of an element is the amount of protons and neutrons ◆ Mass number = # of protons + # of neutrons ● To find the number of neutrons, we can do the “mass number – the atomic number.” This is like doing “protons + neutrons” – “protons,” giving us the amount of neutrons ◆ An isotope is an atom of an element that has a different amount of neutrons. However, the amount of protons stays the same since the element keeps its identity ● All atoms of an element have the same number of protons ◆ This causes a change in mass ● The isotopes of an element differ in their mass numbers since the number of neutrons changes ◆ Example: Carbon 13 and Carbon 14 are isotopes of Carbon. In Carbon 13, the mass number is 13. In Carbon 14, the mass number is 14. To find the number of neutrons in an isotope, we do “mass number – atomic number” (the same formula as above) ◆ In Carbon 13, the mass number is 13. The atomic number is 6 since carbon is the sixth element on the periodic table ● Thus, there are 7 neutrons in Carbon 13 since 13 – 6 = 7. The number being pointed to in this picture is the atomic mass. As you can see, this number is a decimal. When this number is rounded to the nearest whole number, you will get the mass number (protons + neutrons) ➔ This is because the mass number is always a whole number since it is the protons + neutrons, and you can only have a whole number of these subatomic particles ➔ The atomic mass takes the isotopes of the elements and their abundances into consideration, which is why it is a decimal ◆ To learn more about this, read the “Chemistry” study guide Carbon has three naturally occurring isotopes. Carbon-12 is stable, while Carbon-14 is radioactive. ➔ In an atom, the number of protons is equal to the number of electrons. ➔ When changing the number of protons, the atom and element is changed ➔ When changing the number of electrons, the charge (positive or negative) changes ◆ Charged atoms are called ions ◆ A positive ion has more protons than electrons (since protons are positive) ● A positive ion may form if electrons are lost ◆ A negative ion has more electrons than protons (since electrons are negative) ● A negative ion may form if electrons are gained ➔ The periodic table is formatted purposely. It shows many trends and patterns ◆ The columns of the periodic table are called groups or families ◆ The rows of the periodic table are called periods As shown in the image, there are 18 groups (columns) and 7 periods (rows) on the periodic table ➔ We can tell the number of valence electrons an atom has based on the periodic table. ◆ The group number an element is in represents the number of valence electrons an atom of that element has ● Elements in group one have one valence electron ● Elements in group seven have seven valence electrons Note: There is no pattern for determining the number of valence electrons of elements in the yellow space in the image. Those are the transition metals ➔ The Periodic Table can also be used to determine if an element is a metal, nonmetal, or metalloid ◆ First, let’s go over the definitions of metals, nonmetals, and metalloids ◆ Metals are shiny and solid at room temperature. They conduct heat and electricity, are ductile, and are malleable ◆ Nonmetals are mostly gases and brittle solids. They aren’t malleable and ductile, and most do not conduct electricity or heat ◆ Metalloids contain physical and chemical properties of both metals and nonmetals ● They are considered semiconductors, which is important in computer technology ➔ The elements on the left/middle side of the periodic table are metals ➔ The elements on the right side of the periodic table are nonmetals ➔ The few metalloids are between the metals and nonmetals ➔ Note: Although Hydrogen is on the left side of the periodic table, it is not a metal; it is a nonmetal ➔ The groups of the periodic table have specific names, and each group has specific/similar properties ➔ Alkali Metals ◆ Have one valence electron ◆ Most reactive metals; can burst into flames when exposed to air ◆ Reactivity decreases up the group ◆ All elements of group one except Hydrogen ◆ Shiny and white with low melting and boiling points ➔ Alkali Earth Metals ◆ Have two valence electrons ◆ Slightly less reactive than group one metals (Second most reactive) ◆ Harder and denser than Alkali Metals ◆ Shiny with low melting and boiling points ➔ Transition Metals ◆ Groups 3-12 (middle of the Periodic Table) ◆ There are around 40 recognized transition metals ◆ High melting and boiling points ◆ Malleable and shiny ◆ Good conductors of heat and electricity ◆ 2 or 3 valence electrons ◆ High densities ➔ Post-Transition Metals ◆ Parts of Groups 13-16 ◆ Between transition metals and metalloids ◆ Soft and conduct more poorly than transition metals ◆ Melting points are lower than transition metals ◆ Malleable and ductile ◆ Slightly lower densities compared to Transition Metals ◆ The number of valence electrons ranges from 3-6 ➔ Metalloids ◆ Do not fit into metal or nonmetal groups; share properties with both ◆ 3-6 valence electrons ◆ Solid at room temperature ◆ Easily broken and considered brittle ◆ Not as good conductors of electricity as metals, and they aren’t as ductile as metals ◆ Lie between parts of Groups 13 and 16 ➔ Nonmetals ◆ Top right of metalloids ◆ Groups 14-16 (and Hydrogen from Group 1) ◆ Poor conductors of heat and electricity ◆ Exist as gases or solids ➔ Halogens ◆ 7 valence electrons ◆ Highly reactive; only need one electron to complete their outer shell ◆ Nonmetals and fairly toxic ◆ Low melting and boiling points ◆ Poor conductors of heat and electricity ➔ Noble Gases ◆ 8 valence electrons ◆ Complete outermost shells, causing them to be stable ◆ The least reactive group on the periodic table ◆ Usually do not mix or react with other elements ◆ Low melting and boiling points ◆ Colorless and odorless ➔ Lanthanides ◆ Part of the sixth row ◆ Highly reactive ◆ Soft, silvery, and bright metals ◆ High melting and boiling points ◆ Tarnish when exposed to air ◆ Generally aren’t as radioactive as actinides ➔ Actinides ◆ Part of the seventh row ◆ Highly reactive ◆ Similar to lanthanides: soft, silvery, malleable, and ductile ◆ High melting and boiling points ◆ Ignite easily in the air ◆ Generally are more radioactive than lanthanides ➔ Molecules vs Compounds ◆ Molecules are two or more atoms bonded together ● Examples: O2. P4 ◆ Compounds are two or more elements bonded together ● Have fixed ratios and constant compositions ● Examples: NaCl, H2O ➔ Types of Bonding ◆ Ionic Bonds ● Form between a metal (positive charge) and nonmetal (negative charge) ● An atom transfers one or more of its valence electrons to another atom ○ Electrons are transferred, forming ions ● Metals donate their valence electrons to nonmetals since nonmetals are close to filling their valence shells and becoming stable ● These ions with opposite charges will attract, forming an ionic bond ● In sodium chloride (NaCl), sodium (metal) has one valence electron, and chlorine (nonmetal) has seven valence electrons. ○ Sodium donates its one valence electron to chlorine, causing chlorine to have a full valence shell with eight valence electrons ○ Sodium becomes a positive ion because it loses an electron, while chlorine becomes a negative chloride ion since it gains an electron ○ These two oppositely charged ions attract and form the ionic compound sodium chloride ● An electrostatic force holds the oppositely charged ions together to form the ionic bond ● The overall charge of the ionic compound is zero since the charges balance ● Properties of Ionic Bonds ○ Forms an ionic crystal (crystal lattice) ○ High melting and boiling points ○ Conduct electricity when liquids ◆ The ions can separate in water or as liquid ◆ Covalent Bonds ● Forms between two nonmetals ● Electrons are shared between the nonmetals ○ There is no transfer of electrons and there are no ions ○ This is because nonmetals do not donate electrons easily. Instead, they attract other electrons ● The two nonmetals in a covalent bond are both close to having full valence shells, so they will share one or more electrons to achieve this ● The shared electrons will travel around the nuclei of both atoms in the covalent bond ● Water (H2O) is an example of a covalent bond ○ The two hydrogen atoms each share an electron with the oxygen atom ○ This allows the Hydrogens to have full valence shells since they will each have two electrons ◆ Two electrons allow Hydrogen to have a full valence shell since a Hydrogen atom only has one valence shell. The first valence shell is full after having two electrons ○ The Oxygen has a full valence shell as well since it has eight electrons ◆ Note: oxygen has two valence shells, so the outer shell (the valence shell) needs eight electrons to be full ● Properties of Covalent Bonds ○ Low melting and boiling points ○ Poor conductors of electricity ○ Less structured solids (very brittle) ◆ Metallic Bonds ● Form between two metals ● Metals donate electrons easily since their valence shell is low on valence electrons ● The donated electrons in metallic bonds float around the metals ○ These electrons are called delocalized electrons because they move freely ● In the electron sea model, all metals donate valence electrons to form a sea of delocalized electrons ● Properties of Metallic Bonds ○ Moderately high melting and boiling points ○ Malleable (can be hammered out of shape) ○ Ductile (can be drawn in a wire) ○ Durability ○ Good conductors of electricity