Chapter 7: Ionic Bonding • Background (chapter 4): The 3 major subatomic particles: protons, neutrons, and electrons. • The identity of an element is determined by the number ofprotons, (the positively-charged particles) that are found in each atoms' nucleus . • The number of protons is referred to as the "atomic number" of the atom. To make the math easy for us, we say that each proton has a charge = __+1__.
The element hydrogen is the simplest element, and has the simplest atoms: An atom of the element hydrogen contains only _1__ proton in its nucleus; hydrogen's atomic number is 1. • An atom of the element helium contains __2__ protons in its nucleus; helium's atomic number is 2. Each element from atomic number 3 through 118 has a consecutive increase in the number of its protons by ___1_____.
The nucleus is the high density part of an atom, taking up very little of the atom’s_volume_; but, accounting for practically all of its _mass_. This is because, in addition to all the protons, the nucleus also contains neutrons (_neutral_ particles having no electrical charge); and, both the protons and the neutrons are each about 2000 x _more_ massive than the very tiny electron.
While the nucleus by itself is positively-charged (due the protons within it), atoms as a whole are electrically _neutral_ due to their having a number of _negatively_-charged electrons _equal_ to their number of protons. We say that each electron has a charge = -1. The charge of each 1 electron exactly cancels the charge of each 1 proton.
As long as an atom remains neutral, the number of its electrons will equal the number of its protons. • To be neutral means that, although there may be lots of positively– charged and lots of negatively– charged particles present in each atom, overall the charges all cancel out, thus the “net (overall)” charge of a neutral atom equals 0.
Background, Chapter 5: • The electron is the smallest of the 3 major subatomic particles; and, these very energetic particles, though _strongly attracted_ to the positively-charged nucleus, are able to travel quite far away from the nucleus, flying around the nucleus at a speed close to the _speed_ _of_ _light_. • If the nucleus were the size of a marble, the electrons would occupy about the volume of a football stadium. Due to the large volume of space the tiny electrons fly through, atoms are actually mostly _empty__space_. • Niels Bohr suggested that electrons “orbit” the nucleus of an atom at exact distances away from the nucleus.
New information, Chapter 5: • Electrons do NOT_ reside in specific orbits! There is no such thing as an “orbit” (no such thing as an exact distance away from the nucleus) to which electrons are limited! • What _Niels_ _Bohr_ was correct about was this: Electrons have certain “allowable energies_”. It’s notthat they are found in certain allowable orbits away from the nucleus; instead, they are found with certain allowable energies, and they are always outside the nucleus.
When we refer to an electron residing in an _energy_ _level_ (later on in the notes), we are really referring to an amount of energy that an electron has at that moment of time. • And, the electron can’t just have any ole’ amount of energy….The atom itself seems to dictate specific allowable quantities of energy that its electrons might have. • In much the same way, you may have a friend that is always either asleep in class (low energy state), or driving every one crazy with his/her exuberance (high energy state); but, you’ve never seen your friend exhibit a moderate amount of energy.
It’s like that with the electrons….They jump from energy level to energy level, showing sometimes dramatic increases in energy and sometimes showing dramatic decreases in energy; but, never showing an amount of energy, or a change in energy, placing them “ in between” their routinely observed energies. • We believe that electrons, based on the atom that they’re found in, have certain “allowable” energies: This is what is meant by the statement, “The energy of electrons is quantized”. It comes and goes in _packets_ called _quanta_.
When all of the electrons in an atom are at the lowest possible energy state, we say that the atom is at _ground_ _state__ conditions. • When even just one of the electrons in an atom is at a higher energy state (than the lowest possible), we say that the atom is in the _excited__ state_.
Since there is no such thing as an “orbit”, we can’t ever pinpoint a specific location that a specific electron will be found in at a _specific_ moment in time; instead, we can only refer to there being a 90% or better_probability_ of finding an electron in a given region of space at any time. • In much the same way, I can’t predict exactly where I will be at 6:30 am, as I drive in to school; but, I can say that there is a 90% probability I will be somewhere along route 81.
During the early days of the developing “quantum theory of electrons”, De Broglie and Schrodinger called this “probable region”, that is, the region where there is a > 90% chance of finding an electron, an “atomic_ _orbital_". • Atomic orbitals can be envisioned as “electron _clouds_” surrounding the atom's nucleus, in that, like a regular cloud, the edges of any given atomic orbital are _NOT_ distinct boundaries.
Atomic orbitals don’t have defined edges; and, you might have anticipated, there is a 10% chance that an electron will be found outside of its most probable region. • In much the same way, there is a small probability that I might be found somewhere along the shoulder of route 81, or even still, on route 114, at 6:30 am.
Atomic orbitals, like clouds, form different _shapes_. • Some atomic orbitals look like _spheres_, others look like _dumbbells_ (hand weights), and some look like four-leafed_ clovers. • The spherically-shaped orbitals are called "_s___ _orbitals__", • the dumb-bell shaped orbitals are called "__p__ _orbitals_", • the four-leafed clover-shaped orbitals are called "__d__ orbitals_". • The last group of orbitals have complex-shapes and are called "__f__ _orbitals_".
Electrons found in different types of orbitals (s, p, d, or f) will have _different_ amounts of energy. • And, all of the electrons found in the same type of orbital will have about the _same_ amount of _energy_ AS LONG AS _the electrons are in the same energy level.
Niels Bohr was the person who introduced the concept of the “_energy__ __level_”. • Think of an energy level as being similar to one layer of growth in an onion, or one layer of growth in a tree. • Onions and trees are 3–dimensional objects which grow in layers that extend outwards: • In a similar manner, energy levels are 3 dimensional _spaces___ or regions extending outwards from a nucleus. That’s the end of the similarity; but, hold onto that thought.
Each energy level holds electrons which have a range of similar energies. • Electrons having a given amount of energy will be found in a given energy level; and, as electrons gain energy, they will move up to “_higher_” energy levels which are found out _farther away_ from the atom’s nucleus.
Most energy levels (in fact, all but one of the energy levels in an atom) have, during the establishment of _quantum theory__, been sub–divided into what are called _sublevels__, based on the type of orbital(s) found in that energy level. • “s” sublevels contain electrons in “s” orbitals. • “p” sublevels contain electrons in “p” orbitals.
As long as electrons are in the _same_ energy level, the electrons residing in the s-orbitals (the s sublevel) have the _same approximate_ energy. • More energetic electrons (in that same energy level) will be found in the p-orbitals (the p sublevel), even more energetic electrons (in that same energy level) will be found in the d-orbitals (the d sublevel), and the __most__ energetic electrons (in that same energy level) will be found in f-orbitals (the f sublevel).
Memory aid for the order, from least energy to most energy within an energy level: spdf ... __some people don’t fart_
In class we will need to become comfortable working with sublevels; and, we will find out that the two most important sublevels (when trying to understand_chemical_ changes) are the __s_ sublevel and the _p_ sublevel. • The symbol, “1s”, is our symbol for the “1 s _sublevel_”, and is our way of saying, “first__ energy level”, “s orbitals only”. • The symbol, “2p”, is our symbol for the “2 p sublevel”, and is our way of saying 2nd energy level, ___p___ _orbitals__ only. • Etc. What is our symbol for saying, “3rd energy level, d orbitals only”? ___3d__
There are specific numbers of electrons that an orbital can hold maximally, specific numbers of electrons that sublevels can hold maximally, and specific numbers of electrons that energy levels can hold maximally. • Thus, there is a specific organization of orbitals, sublevels, and energy levels, within each atom. • There are several different patterns that occur. • Understanding these patterns will help you keep the organization straight in your mind.
– Here is a pattern between energy levels and sublevels. Can you see the pattern? • 1st energy level has ____1_____ sublevel, identified as: __1s__ • 2nd energy level has _2__sublevels, • identified as: __2s__ __2p__ • 3rd energy level has __3__ sublevels, • identified as: __3s__ _3p___ _3d___ • 4th energy level has _4_ sublevels, • identified as: _4s___ __4p__ _4d___ _4f___ • 5th energy level has _4_ sublevels, • identified as: __5s__ _5p___ _5d___ _5f___ • 6th energy level has ___3__ sublevels, • identified as: _6s__ __6p__ _6d___ • 7th energy level has __2__ sublevels, • identified as: _7s__ __7p__
There are also patterns in the number of orbitals in a sublevel, and in the maximal number of electrons in a sublevel. Can you see the patterns? = X
By combining the two patterns, you can see that the formula for determining the number of electrons held maximally within the first four energy levels = __2n2___ • where “n” = _number of the energy level (1 through 7) 118 118
The “__diagonal_ rule” (next page) uses the sublevel pattern to help scientists generate what is called an “_electron_ _configuration_” and an _orbital_ _diagram_ when the atom is in its _ground state_ (lowest energy) condition. • You are always allowed to use your reference packet, which includes a diagram of the diagonal rule (and the periodic table, of course, is always available); • but, while the energy ladder is displayed in our classroom – it is not guaranteed to be visible to you during the unit exam, the midterm and/or the final exams.
To use the diagonal rule, just read off the sublevels one at a time, as you follow along the slanted arrows (from _right_ to _left_), starting at the top of the diagram and ending at the bottom. Additionally, state the number of electrons found in each sublevel.
s-2e p-6e d-10e f-14e • 1s • 2s 2p • 3s 3p 3d • 4s 4p 4d 4f • 5s 5p 5d 5f • 6s 6p 6d • 7s 7p
When reading from the diagonal rule correctly, you generate a correctly ordered list of sublevels: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p When working with an atom of neon, which only has 10 electrons, its electron configuration would be this: 1s2 2s2 2p6 *as you follow the diagonal arrows, add up the exponents till you get to 10, then stop. See note below. In other words, an electron configuration is a list of sublevels, each being “filled” with the maximum number of electrons for that given sublevel. The list of maximum electrons is given to the right side of the diagonal rule diagram above. What would be the electron configuration for an atom of argon? Atomic number = 18. so, 18 electrons. 1s2 2s2 2p6 3s2 3p6 2 + 2 + 6 + 2 + 6 = 18 Follow the diagonal arrows, Maximum exponent for every “s” = 2. Max every “p” = 6. max “d” = 10. max “f” = 14”
Neon, 10 electrons: Read directions (next page in notes) Try Argon, 18 electrons on your own:
Lastly, as per the instructions on your last page of notes, try writing out an electron configuration and orbital diagram for nitrogen. • If you don’t yet have the webquest handout, now is the time to request one!