Science Periodic Table PDF

CHEMICAL ELEMENTS Early attempts in classifying elements - During the early 1800s, there were only about 30 elements known to scientists - By the turn of the new century, the number of elements known doubled and this resulted in the efforts of classifying and organising the elements JO...

CHEMICAL ELEMENTS Early attempts in classifying elements - During the early 1800s, there were only about 30 elements known to scientists - By the turn of the new century, the number of elements known doubled and this resulted in the efforts of classifying and organising the elements JOHANN WOLFGANG DÖBERENEIR - In 1829, he observed that several elements that have similar properties could be grouped into three called TRIADS. - In each triad, the the middle element was approximately the average of the first and third element’s properties - The atomic weight is the most common and consistent property that is observed and has a pattern - Add the atomic mass of the first and third element and divide by 2 DOBEREINER’S TRIADS - However, this identification of new elements made this model obsolete and a new scientist named John Newlands created a different model as the newly discovered elements did not fit into the triads JOHN NEWLANDS - In 1865, he presented another way of classifying elements which is called the Law of Octaves - He arranged the elements in order of increasing atomic mass - He noted that in every eight elements, there would be a repetition of similar properties LAW OF OCTAVES - Newlands noticed that this pattern was similar to musical octaves in where it repeats every eighth note - However this method of classifying the now 62 elements was met with a lot of resistance in the scientific community as some elements with dissimilar properties were grouped together LIMITATIONS OF TRIADS AND LAW OF OCTAVES - The elements that were discovered later could not fit into the triad and octave patterns - Even if these primitive methods of classifying elements did not work out, they laid the foundation for the development of the modern periodic table DIMITRI MENDELEEV - He arranged the elements in increasing atomic mass and noticed that its properties had repeating patterns but also noticed that there were gaps in his arrangement - Mendeleev realised that some elements may still be undiscovered with this, he left gaps in his table for the elements that have not yet been discovered - However, there were discrepancies within Mendeleev’s table as the increase in atomic mass was not regular while moving from one element to another. Hence the number of elements yet to be discovered was not predictable MENDELEEV’S PERIODIC TABLE - Despite leaving gaps in his periodic table, some elements still appeared to be positioned incorrectly - For example tellurium and iodine didn’t fit with the properties of the elements they were grouped with HENRY MOSELY (1913) - He observed that frequencies of x-rays emitted from elements could be correlated better w/ atomic number - He then realised that the elements should be arranged in order of increasing atomic number than increasing atomic mass MODERN PERIODIC TABLE - Mosely summarised his findings by stating the modern periodic law “The properties of the elements are periodic functions of their atomic mass” THE PERIODIC TABLE OF ELEMENTS - It organises all discovered chemical elements in rows (called period) and columns (called groups) according to increasing atomic number - Scientists use the periodic table to quickly refer to information about an element like atomic mass and chemical symbol - The periodic table’s arrangement also allows scientists to discern trends in element properties HOW TO USE THE PERIODIC TABLE OF ELEMENTS Groups and periods - Groups of vertical columns of elements -> Groups IA-VIIIA : Denotes similar number of valence electrons (How many electrons in the outermost Shell) -> Groups 1-18 : Denotes similar chemical properties; recommended by IUPAC - Periods or horizontal rows of elements -> Each row corresponds to the number of energy level(s) that can hold electron(s) -> How many circles ALKALI METALS - Elements that form alkaline solutions w/ water - Alkaline solutions are strong bases that can neutralise acids - 1 valence electrons ALKALI-EARTH METALS - Elements that form alkaline solutions w/ water - Fire-resistant substances - Reactive w/ other elements - 2 valence electrons TRANSITION METALS - Harder than alkali metals - Less reactive w/ water - Used for structural purposes (buildings, tables, etc.) LANTHANIDES - Inner transition metals - Shares common properties with the first element : Lanthanum (La) - Difficult to purify ACTINIDES - Inner transition metals - Shares common properties with the first element : Actinium (Ac) - Difficult to purify POST-TRANSITION METALS - Located between transition metals and metalloids - Melting and boiling points generally lower than transition metals - Semiconductors, some form of alloys METALLOIDS - Chemical elements whose physical and chemical properties fall in between metal and nonmetal categories REACTIVE NON-METALS - Non-metals are the elements which form negative ions by accepting or gaining electrons. Non-metals usually have 4, 5, 6 or 7 electrons in their outermost shell. - They usually gain electrons when reacting with other compounds, forming covalent bonds - Mostly gases, some solids and liquids HALOGENS - Salt forming non metals - Have some seven valence electrons; because halogens have one electron missing - They form negative ions and are highly reactive NOBLE GASES - Has 8 valence electrons - All unreactive gases that tend not to combine w/ other elements PERIODIC TRENDS - Periodic trends are patterns that are present in a periodic table that show the different aspects of a certain element, including its radius and other properties MAJOR ATOMIC TRENDS - Atomic radius - Ionic radius - Ionisation energy  - Electron affinity - Metallic character - Electronegativity ATOMIC RADIUS - Helps us understand why some molecules fit together and why other molecules have parts that get too crowded under certain conditions - The size of an atom is defined by the edge of its orbital. However the orbital boundaries are fuzzy, and variable under different conditions. In order to standardise the measurement of atomic radii, the distance between the nuclei of two identical atoms bonded together is measured - Atomic radius is defined as one half of the distance between the nuclei of identical atoms that are bonded together - UNIT OF MEASUREMENT : Picometer (pm) or 10^(-12)m WHAT AFFECTS THE SIZE IN AN ELEMENT? Number of electron shells : More shells mean a larger atomic radius because the outermost electrons are farther from the nucleus Nuclear charge : As the number of protons in the nucleus increases (based on the atomic number), the greater positive charge pulls electrons closer, reducing the atomic radius. This is why atomic radius generally decreases across a period (left to right) on the periodic table Shielding effect : Inner electrons can shield the outer electrons from the full attraction of the nucleus. More shielding reduces the pull on the outer electrons, increasing the atomic radius *The more protons an element has, the smaller it is as the attraction of positive and negative is stronger *The more shells an element has, the bigger it is as the shells are added increasing the size (The size of an element is determined by the valence electrons) GROUP TRENDS IN ATOMIC RADIUS - The atomic radius of atoms generally increases from top to bottom within a group. As the atomic radius increases down a group, there is again an increase in the positive nuclear charge (proton) - There are some small exceptions, such as the oxygen radius being slightly greater than the orbitals from lower energy levels - There is also an increase in the number of occupied principal energy levels. Higher principal energy consists of orbitals which are larger in size than the orbitals from lower energy levels - The effect of the greater number of principal energy levels outweighs the increase in nuclear charge, and so atomic radius increases down a group PERIOD TRENDS in ATOMIC RADIUS - The atomic radius generally decreases from left to right across a period. - Within a period, protons are added to the nucleus as electrons are being added to the same principal energy  - These electrons are gradually pulled closer to the nucleus because of its increased positive charge. Since the force of attraction between nuclei and electrons increase, the size of an atom decreases - The innermost shell can only hold a maximum of 2 electrons PERIODIC TRENDS in IONIC RADIUS and IONIZATION ENERGY - When atoms have a fewer than eight valence electrons, they tend to react to other atoms and form more stable compounds - When atoms gain and lose electrons, they form ions - Ions are charged atoms (+/-) OCTET RULE (outer shell) - The octet rule refers to the tendency of atoms to prefer to have 8 electrons in the valence shell. When atoms have fewer than 8 electrons, they tend to react and form more stable compounds - To be stable, an atom will gain, lose, or share electrons to compete the outermost energy level (electron shell) - Metals in group 1A, 2A, and 3A will form ions with 1+, 2+ and 3+ charges respectively this means that they have an extra ion (based on their number) and they have to remove it to fit the 8 electron rule or octet rule *1A = 1+ *2A = 2+ *3A = 3+ - Non metals in group 5A, 6A, and 7A will form ions with 1-, 2- and 3- charges respectively this means that they need a certain amount of ions (subtract from their group) and they have to add electrons to fit the 8 electron rule or octet rule *5A = 3- *6A = 2- *7A = 1- - Noble gases or elements in group VIIIA (8) do not form any ions at all IONIZATION ENERGY - This is the minimum energy in (kj/mol) required to remove an electron from a gaseous atom in its ground state - The magnitude of ionization energy is a measure of how tightly the electron is held in the atom - The higher the ionization energy, the more difficult it is to remove the electron A small atom requires high ionization energy because the closer an electron is to the nucleus, the greater the nuclear attraction. This is due to the increase in the nuclear charge resulting in the stronger forces of attraction exerted on the electrons IONIZATION - The process in which an atom loses or gains an electron to form an ion TRENDS IN IONIZATION ENERGY (Talks about how much energy is used to produce an ion) - Group trend ~ The first ionization energy generally increase as you move down a group. This is due to the increase in the atomic size of elements down a group - Period trend ~ For the representative elements, the first ionization energy generally decreases from left to right across a period *Ions are atoms that carry a (+) or (-) charge IONIC RADIUS - Cations have smaller radius than the neutral atoms from which they are formed (Cations - Positive) ex. 2+, 3+, etc. - Anions have larger radius than the neutral atoms from which they are formed (Anions - Negative) ex. 2-, 3-, etc. Cations are smaller because it’s formed by the loss of electrons to fit the octet rule. We remove the negative (electron) making it positive. This decreases the size as one electron is like one shell. Anions are bigger because we have to add an electron to fit the octet rule. We remove the positive (proton) making it negative. This increases the size as adding one electron is like adding one more layer or shell. - Neutral atoms are in the middle of Anion and Cation If an element removes electrons to achieve the octet rule, it is a positive. If an element needs electrons to achieve octet rule, it is a negative. Ionic radius increasing going down a group (horizontal) Ionic radius decreasing across a period left to right (vertical) Ionisation energy decreasing down a group (horizontal Ionisation energy increasing a period from left to right (vertical) PERIODIC TRENDS : ELECTRONEGATIVITY (How much they want to attract electrons for octet rule) - Atoms form a chemical bond to become stable. With these chemical bonds, atoms of some elements have a greater ability to attract the valence electrons involved in the bond than other elements Non polar covalent bond - Bonding electrons shared equally between two atoms. No charges on atoms Polar covalent bond - electrons are shared unequally between two atoms. Partial charges on atoms Ionic bond - One is giving and one is receiving (Cation and Anion) What is electronegativity? - It is a measure of how much an atom can attract a bonding pair of electrons - How much an atom wants to lose or gain an electron CATIONS : REMOVE ELECTRONS ANIONS : RECEIVE ELECTRONS Hydrogen bonds w/ chlorine electronegative - Chlorine attracts the electron pair rather more than hydrogen does - The Chlorine end of the bond has more than its fair share of electron density and so becomes slightly negative. At the same time, the hydrogen end becomes slightly positive Same electronegativity values are bonded - Bonding electrons are shared equally between the two atoms - Both atoms of chlorine are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so it will be found on average half-way between the two atoms Bond metal + non metal - If an atom like chlorine is a lot more electronegative than an atom like sodium, then the electron pair is dragged right over to the chlorine end of the bond - Sodium has lost control of its electrons, and chlorine has complete control over both electrons. Ions have been formed TRENDS IN ELECTRONEGATIVITY - Decreases down a group because the increased number of energy levels puts the outer electrons very far away from the pull of the nucleus - Increases across a period because the number of charges on the nucleus increases. This attracts the bonding pair of electrons more strongly ELECTRONEGATIVITY : Electronegativity refers to the amount of energy an atom wants to gain or lose an electron. The higher Electronegativity an element has, the more it wants to gain an electron to achieve the octet rule to make the atom stable. If an atom has 4 or more electrons in their valence shell, the stronger the Electronegativity energy so that it can attract more electrons so that the valence shell will have 8 electrons. The lower Electronegativity an element has, the more it wants to lose the electron/s in its valence shell. The reason for this is that the second to the last shell already has 8 valence electrons so the atom will lose the electrons in the valence shell so that the second to the last shell will replace the valence shell and achieve stability and octet rule as the “valence shell” has 8 electrons. ELECTRON AFFINITY - Defined as the charge in energy (in kj/mol) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion. In other words, the neutral atom’s likelihood of gaining an electron - An increase in electron affinity corresponds to an increase in the tendency to accept electrons *NOBLE GASES HAVE 0 electron affinity, Electronegativity and ionisation energy because they are already stable meaning they don't want electrons anymore* WHY NON METALS HAVE A GREATER ELECTRON AFFINITY THAN METAL ATOMS?  - Non metals have a greater electron affinity than metals because their atomic structure allows them to gain electrons rather than lose them GROUP : Generally decreases as you go down the group 1. Electrons are placed in energy levels further away from the nucleus 2. The atom ‘does not want’ to gain electrons because there is minimal charge on the outer energy levels from the nucleus 3. The shielding effect increases, causing repulsion between the electrons, thus they move further from each other and the nucleus itself IRREGULARITIES IN THE TREND IN ELECTRON AFFINITY - The electron affinity of a group 2A element is lower than that of a corresponding group 1A element - The electron affinity of a group 5A element is lower than that for a corresponding element group 4A element - The electron affinity for fluorine is less than that for chlorine *These irregularities or exceptions are due to the valence electron configuration of the elements involved ATOMIC ORBITALS - Electrons spend most of their time in an electron shell down into one or more sub shells, which are simply sets of one or more orbitals - Subshells are designated by the letters s, p, d and f & each letter indicates a different shape